
Class & X1 -S-^ 
Book ,Mll 3 



tOF-OUGHT DEPOSIT, 




INTRODUCTION 

TO 

GENERAL CHEMISTRY 



McCOY and TERRY 




/ 



INTRODUCTION 

TO 

GENERAL CHEMISTRY 



By 

HERBERT N. McCOY 

and 

ETHEL M. TERRY 



CHICAGO, ILLINOIS 
1919 






vA 



\S 



Copyright 1919 By 
Herbert N. McCoy and Ethel M. Terry 



All Rights Reserved 



Printed May 191 9 



MAY k6 I 



Composed and Printed By 

The University of Chicago Press 

Chicago, Illinois, U.S.A. 

©CLA525660 



V~vA 



CHAPTER I 



INTRODUCTION— LAWS OF GASES 



i. A Knowledge of Physics Prerequisite for Chemistry. — 

The sciences of physics and chemistry are so closely related that 
the latter may be considered an extension of the former. A 
knowledge of physics is therefore necessary for an adequate 
understanding of chemistry, and it is to be assumed that the 
student taking up chemistry has had at least a one-year high- 
school course in physics. 

2. The Three Forms of Matter: Gases. — In his work in 
physics, the student will have learned the meaning of the term 
matter, which may be defined as anything which occupies space 
and has weight. He will have learned, also, that matter may 
exist in three forms: solid, liquid, and gaseous. Since gases are 
less tangible than solids and liquids, we shall first take up the 
study of air, the most familiar of all gases. That 

air has the two attributes just mentioned as belong- 
ing to all forms of matter may readily be shown by 
experiment. 

3. Air Occupies Space and Has Weight. — If a 
drinking glass or beaker be thrust, mouth down- 
ward, into a vessel of water, the water does not 
enter until the glass is tilted to allow the air to 
escape. This shows that air occupies space. 

That air has weight may be shown by weighing 
a flask, -first empty and afterward filled with air. 
The flask (Fig. 1) should be round-bottomed and 
have a capacity of 250 to 500 c.c. It is fitted with 
a tight rubber stopper carrying a glass stopcock. The air is 
first pumped out by means of an efficient air pump; the stop- 
cock is then closed and the flask counterbalanced with weights. 
When the stopcock is opened the inrush of air can be heard, 
and it is easy to observe that there is an appreciable increase 




Fig. i 



2 Introduction to General Chemistry 

in weight. Since air occupies space and has weight, it is un- 
doubtedly a form of matter. 

One liter of air weighs more than a gram and the air contained 
in a room 12 feet square and 12 feet high would weigh about 100 
pounds. At the earth's surface air exerts a pressure of about 
15 pounds on every square inch of surface. The existence of 
this pressure may readily be shown by means of the following 
experiment. A tin can with a narrow neck (such as is often used 
for shipping alcohol, etc.) and of about 1 gallon capacity is 
fitted with a stopper carrying a glass tube, by means of which 
the air filling the can may be pumped out. Usually, before the 
exhaustion of the air is complete, the can is crushed by the 
pressure of the air on the outside — a pressure which is now 
no longer balanced by the equal and opposite pressure on the 
inside. 

4. The Effect of Pressure on Volume: Boyle's Law. — The 
atmospheric pressure is measured by means of the barometer. 
At the sea-level the normal barometric pressure serves to support 
a column of mercury 76 cm. high. The effect of pressure upon 
the volume of air was first studied by Robert Boyle in the seven- 
teenth century. Boyle found that the volume of a given portion 
of air was inversely proportional to the pressure. This relation 
is known as Boyle's law. If we represent the pressure by P and 
the volume by V, then PF = a constant. 

5. The Effect of Temperature on Volume: The Law of 
Charles. — In scientific work we use the Centigrade thermometer, 
the scale of which is so constructed that the freezing-point of 
water is o°, while the boiling-point is ioo°. The effect of tempera- 
ture upon the volume of a given portion of air at a fixed pressure 
was studied over a century ago by Charles and by GayLussac. 
It was found that the volume of the air increased 1/273 of its 
volume at zero for each increase of i°C. This statement is 
known as the law of Charles, or sometimes also as the law of 
Gay Lussac. 

6. The Gas Thermometer: Absolute Temperature. — An 
experiment will show that if 273 c.c. of air contained in a flask 
or cylinder at o° C. is heated to ioo° C. the volume will change to 



Introduction — Laws of Gases 3 

373 c.c. Such an apparatus is called an air thermometer and 
temperatures may be measured in this way instead of by 
the expansion of mercury, as in ordinary thermometers. At 
i° C. the volume of the air is 274 c.c; at 2 it equals 275 c.c. 
and thus the volumes in the following table correspond to 
the temperatures given. 



TABLE I 



Volume 


Degrees 


Volume 


Degrees 
Centigrade 


in c.c. 


Centigrade 


in c.c. 


373 minus 2 


73 = 100 


293 minus 273= 20 


372 " 


' = 99 


283 


= 10 


363 " 


< = 90 


273 


" = 


323 " 


' = 5o 


263 


1 " =-10 


313 


' = 40 


253 


' " =-20 


303 " 


1 = 30 


243 


" =-30 



Since the zero of the Centigrade thermometer is arbitrarily 
chosen, being the temperature of the freezing of water (the stu- 
dent is already familiar with the Fahrenheit zero, which is at a 
lower temperature), it would be possible and convenient to use 
a temperature scale in which the volumes of the air in the air 
thermometer as described are taken as the temperatures. Since 
temperatures on the Centigrade scale are obtained by subtracting 
273 from the corresponding air thermometer temperatures, the 
zero of the air thermometer or gas scale must be 273 degrees lower 
than the Centigrade zero, or 273 degrees below the freezing-point. 
These air-thermometer temperatures are usually called the abso- 
lute temperatures; the absolute temperature may therefore be 
defined as the Centigrade temperature plus 273 degrees. Since 
most other gases act like air they may be used in the gas ther- 
mometer, and it is evident that if a certain amount of gas is used 
in such an experiment, no matter what its volume may be at the 
freezing-point of water, the volume will always vary with the 
temperature in the same ratio as the absolute temperature, pro- 
vided the pressure on the gas is kept constant. 

7. Problems. — We may now consider a few simple problems 
based on the two laws of gases just discussed. 



4 Introduction to General Chemistry 

Problem i: The volume of a certain amount of air at 27 C. 
is 1,000 c.c. What would its volume be at 12 7 C. if the pressure 
is kept constant? 

Centigrade temperature-}- 2 73 = absolute temperature 

2 7 +273 := 3 00 
i27°+273=4oo 

The volume of the gas must therefore increase in the ratio of 
400 to 300, or it will become 

1000 c.c.X — = 1333 -3 c.c. 

Problem 2: Let the original pressure on the gas in Problem 1 1 
be 60 cm. of mercury (or — of the ordinary pressure of the atmos- 
phere) . What will be the final volume of the gas if the pressure 
is increased to 100 cm. of mercury? An increase of pressure 
must decrease the volume of the gas, and in the ratio of the 

pressures, 60 to 100, or by — . 

a) Let the change of pressure come after the change of 
temperature as given in Problem 1 : 

then 

1333 .3 c.cX = 800 c.c. (final volume) Ans. 

100 

b) Let the change of pressure take place first : 

1000 c.cX = 600 c.c, volume after the pressure change. 

100 

The temperature change would then change the volume as 
follows : 

60c c.c.X — = 800 c.c. (final volume) Ans. 
300 

It is thus seen that the same answer is obtained, no matter 
which step in the problem is worked first, so the whole problem, 
1 and 2 together, may be stated in one expression as follows : 

400 60 
1000 c.c.X 11 — X — = 800 c.c. Ans. 
300 100 



Introduction — Laws of Gases 5 

Problem 3: Suppose that 1,000 c.c. of air at 20 C. and 70 cm. 
pressure is cooled to o° and that at the same time the pressure 
is increased to 76 cm. Find the final volume. 

When a gas is at the temperature of o° C. and under a pressure 
of 76 cm. (the normal atmospheric pressure at sea-level) it is said 
to be at standard conditions. 

Problem 4: Find the volume at standard conditions of 400 c.c. 
of air measured at 30 and 72 cm. 

8. Steam Is the Gaseous Form of Water. — It is well known 
that when water is heated it passes into steam. The white cloud 
which is frequently spoken of as steam is not really steam, but 
is composed of minute droplets of water. If we boil water in a 
glass flask the space above the water is rilled with steam, but 
we notice that the steam is entirely invisible and that the visible 
cloud forms only when the steam cools and condenses to liquid 
droplets. Water in the form of steam is, like air, a gas. When 
we boil any liquid like alcohol or mercury the liquid passes into 
the state of a gas or vapor, as it is sometimes also called. The 
gas or vapor when cooled condenses to the liquid form of the sub- 
stance. 

9. Change of Form of Matter with Change of Temperature. — 
Just as water when cooled solidifies to ice, so every other liquid 
substance solidifies when sufficiently cooled. We speak of steam 
and ice as the gaseous and solid forms respectively of water. 
The substance known as moth-balls is called naphthalene by the 
chemist; it is a solid at ordinary temperatures, but when heated 
it melts to a colorless liquid, and when heated still hotter it 
boils, giving a colorless vapor, which is naphthalene in the form 
of a gas. When this gas is cooled it condenses to a liquid, which 
when cooled still further solidifies or freezes, giving solid naph- 
thalene again. Behavior like that of water and naphthalene 
is met with in the case of very many other substances. They 
can exist in three different forms, gas, liquid, and solid, according 
to the temperature. 



CHAPTER II 



THE BURNING OF SUBSTANCES— OXYGEN 



Fig. 2 



io. Burning Substances Require Air. — The history of chem- 
istry shows that the discovery of the real nature of the process of 
burning was one of the most important, if not the most important, 
in the development of the whole science. That air is needed for 
the burning of a substance is, in general, well known, and can 

easily be shown by many simple 
experiments. For example, if we 
place an inverted drinking-glass 
over a burning candle standing on 
a table (Fig. 2), the flame quickly 
grows smaller and smaller and soon 
goes out, the glass having cut off 
l the needed supply of air. 

A still more interesting and in- 
structive experiment may be made 
with phosphorus, a substance which burns very readily in the 
air, giving off clouds of white 
smoke. A piece of phosphorus of 
the size of a pea is placed on a 
cork floating on water and covered 
with a bell-jar (Fig. 3). When a 
heated wire passing through the 
tight-fitting stopper of the jar is 
brought in contact with the phos- 
phorus the latter takes fire and 
burns with the production of light 
and heat and the formation of a 
cloud of white smoke. At the 
same time the level of the water 

inside the bell-jar first falls a little and later rises; but while 
there is still a large volume of air left above the water on which 
the cork floats, the flame dies out and the burning ceases. By 



77 



*» 



N 

^ 

O 



s //;; 



Fig. 3 



The Burning of Substances — Oxygen 7 

the time the bell-jar and its content have become cold, the 
cloud has disappeared and the water has risen on the inside so 
that the volume of the remaining air is seen to be about four- 
fifths of the original volume. It follows that about one-fifth 
by volume of the air has disappeared. 

Further examination also shows that much of the phos- 
phorus still remains unburned. Why, then, should the burn- 
ing stop while there is still four-fifths by volume of the air left 
in the jar? The answer to this question may be made when 
we find that, try as we may, we cannot make phosphorus or 
anything else burn in the air remaining in the jar. We therefore 
conclude that the remaining air is different from common air. 
The correctness of this conclusion is supported by the fact that 
small animals, such as mice, suffocate at once if allowed to breath 
this remaining portion of the air. The facts just considered make 
it seem probable that one-fifth of the air is different from the 
balance, and that it is this portion which takes part in the burning 
of substances and which is necessary for the respiration of 
animals. 

Everyday experience would seem to indicate that wood, coal, 
paper, gasoline, etc., are completely destroyed when they are 
burned. Wood and coal leave a small amount of ash when 
burned, but nothing visible remains in the cas'e of gasoline and 
other oils. Since we have found that water in the form of steam 
is invisible, it is possible that the substance burned may have 
passed into an invisible form and thus escaped notice. 

There are many substances which burn very readily and in 
so doing leave behind large amounts of ash; the experimental 
study of the burning of such substances leads to important con- 
clusions. We may now consider two typical cases of this sort. 

11. The Burning of Magnesium. — The metal magnesium, 
which is used in photographic flash lights, will burn very readily 
in air, either in the form of powder or thin ribbon. In either case 
we notice that a white ash is left. If we collect and weigh the ash 
from the burning of a weighed piece of magnesium ribbon we 
find that the ash weighs more than the original metal ribbon. 
The actual experiment is best carried out by placing about one 



8 



Introduction to General Chemistry 



O 



gram of magnesium, in the form of wire (Fig. 4) or ribbon, in a 
porcelain crucible, having a cover, and then weighing crucible and 

contents. The magnesium is then 
ignited and the cover so adjusted 
that some air can enter, but that 
the dense cloud of white snjoke is 
largely held back in the crucible. 
After the burning is finished and 
the crucible has cooled and the 
whole is again weighed, it will be 
found that there has been a consider- 
able increase in weight. 
12. The Burning of Iron. — Iron powder or filings burn readily 
when thrown into a flame, and in a similar manner we find that 
the burned iron or iron ash, as we might possibly call it, is heavier 
than the original metal. In order to show this by experiment, 
we may suspend on one side of a balance (Fig. 5) a horseshoe 



Fig. 4 




Fig. 5 



magnet which has been dipped in iron filings, and counterpoise 
the magnet and adhering iron by adding small shot or sand to 
the other pan of the balance. By the application of a flame, the 
iron, which now presents a large surface to the air, may be 
ignited. As it burns with a dull glow we observe a gradual 
increase in its weight, and, while there is no noticeable change 
in its volume, the cold residue, which we may call iron ash, is 



The Burning of Substances — Oxygen 



seen to have lost its metallic luster and taken on a dead black 
color. We find, thus, that iron ash is heavier than the iron burned. 
If we seek the cause of this increase in weight, we may get a 
hint when we remember that for the burning of a candle air is 
required, and that, moreover, part of the air disappeared when 
phosphorus was burned in it. What, then, becomes of the 
weight of the one-fifth of the air that disappeared? Is it added 
to the weight of the iron, so as to increase the weight of its ash? 
The facts presented in the next paragraph will furnish the 
required answers. 

13. Lavoisier's Experiment with Mercury. — An experiment 
which turned out to be one of the most important made in the 
early 'development of the 
science of chemistry was 
carried out by the great 
French chemist, Lavoisier, 
in the latter part of the 
eighteenth century. The 
arrangement in this classic 
experiment is shown in 
Fig. 6. The retort (the 
glass vessel with the long 
bent neck) was partly filled 
with mercury (quicksil- 
ver) ; the space above the mercury contained ordinary air, which 
also filled the bell-jar with which the neck of the retort communi- 
cated. The bell-jar stood in a shallow vessel containing mercury, 
which served to prevent outside air from passing into or out of 
the jar. The mercury in the retort was now heated by means of 
a charcoal stove for a period of several days. The heating first 
caused an expansion of the air; but as time went on a gradual 
contraction occurred, which entirely ceased after several days, 
whereupon the heating was stopped. The volume of the air left 
in the entire apparatus when brought to its original temperature 
and pressure was practically four-fifths of what it had been at the 
start. The surface of the mercury in the retort was found 
to be covered with a red powder, which may be considered 




Fig. 6 




io Introduction to General Chemistry 

as analogous to the white ash formed in the burning of 
magnesium or the black ash formed by the burning of iron 
filings. 

14. Heating the Red Ash of Mercury. — If we take some of the 
red ash of mercury, place it in a glass test tube, and heat it very 
strongly (Fig. 7), we find that it changes in a remarkable way: 
first it turns black, and then at red heat it gradually grows 
smaller, until after a few minutes none of it remains. At the 
same time, however, on the cooler part of the 
wall of the tube a silvery-looking coating has 
appeared, which when the tube has cooled 
may be brushed to the bottom of the tube, 
and is then readily seen to consist of drops of 
liquid mercury. Thus by heating the red 
powder to a higher temperature than that used 
in its formation, mercury is reproduced. But 
this is only half the story. 

The more important part remains to be 
II told. Lavoisier reasoned about the matter 

somewhat as follows: If burning substances 
require air; if a part of the air disappears (in 
some cases at least) during burning; if in the 
burning of metals like magnesium and iron the ash is heavier than 
the metal burned; if, as is indeed a fact, air has weight; is it not 
possible that the burning substance unites with a part of the air 
to form a new kind of substance, and that this new substance, 
for example, magnesium ash, is heavier than the substance 
burned because it contains not only the latter but also a part of 
the air? Perhaps also the red ash formed by the gentle heating 
of mercury in contact with air is also made up of mercury and 
something taken from the air. Perhaps the one-fifth of the air 
that vanished has combined with the mercury to form the red 
ash. If all these suppositions are true, perhaps when the red 
ash was changed again into mercury by being strongly heated 
there was set free at the same time the part of the air which by 
originally uniting with the mercury produced the red ash. If 
all this were true, how could it be proved? Let us see. 



The Burning of Substances — Oxygen n 

15. The Active Part of the Air: Oxygen. — The part of the air 
which disappeared may be just that part which causes substances 
to burn. If it were to be obtained pure, free from the inert four- 
fifths which does not support burning (combustion), it ought to 
support combustion far better than common air. This is a matter 
easily put to the test of experiment. Let us again heat some 
of the red powder in a test tube and at the same time thrust into 
the tube a burning wood splint. We see that it burns much more 
fiercely and brightly than in common air. Furthermore, if we 
have no flame, but only a tiny spark on the end of the splint, 
we see that when thrust into the tube above the heated red ash 
the spark bursts into a vigorous flame. The suppositions seem 
to be true. Lavoisier was led in this way to the discovery of 
the secret of the nature of burning. He called the gas formed 
by the heating of the red powder oxygen. This gas forms one- 
fifth by volume of the air and is the part of the air which is necessary 
for the burning of substances. The other four-fifths by volume 
of the air is inert; it does not support combustion; neither does 
it support the respiration of animals. Lavoisier called it azote; 
we call it nitrogen. 

16. The Properties of Oxygen. — Oxygen is an invisible gas 
like air; it has no odor and it supports combustion far better than 

does air. By the same method as that employed 
in the case of air (chap, i), we may readily find 
Oo that 1 liter of oxygen at a temperature of o° and 

76 cm. pressure weighs 1.43 g. It has, there- 
H_l fore, a somewhat greater density than air, of 

which 1 liter weighs 1.29 g. Further evi- 
dence that the explanation of the nature of 
burning, given in the preceding paragraph, is the 
correct one is furnished by experiments which 
we may now consider. 

17. Burning Iron in Oxygen. — If we place a 

gram or two of iron filings and a minute piece of 

phosphorus on a piece of asbestos paper in the 

bottom of a 300-c.c. round-bottomed flask filled with pure 

oxygen and fitted with a rubber stopper and a glass stopcock 




12 Introduction to General Chemistry 

(Fig. 8), we shall find that the weight of the flask with its con- 
tents does not change if by heating we cause the iron to burn 
in the oxygen. Now, we know that when iron burns, the prod- 
uct weighs more than the original iron. We know also that 
oxygen has weight and that the total weight of the flask with its 
contents has not changed during the burning. What then is the 
cause of the increase of the weight of iron when burned ? If we 
open the stopcock while its open end is held under water, we find 
that the water nearly fills the flask. We must conclude that the 
oxygen has disappeared. Is it not reasonable to suppose that 
the ash resulting from the burning of iron is composed of the iron 
originally taken and the oxygen which has disappeared? Our 
experiment has shown that the weight of this ash is precisely the 
same as the combined weights of the iron and the oxygen which 
disappeared in the burning. 

It will readily be seen that the experiment with iron is similar 
to that made by Lavoisier with mercury — with the difference 
that iron burns rapidly, whereas mercury changes but slowly in 
oxygen. Furthermore, the fact that the red ash of mercury 
when strongly heated gives again mercury and oxygen makes 
it practically certain that the red ash was formed by the combina- 
tion or union of mercury with oxygen which composed part of 
the original air used in Lavoisier's experiment. Instead of iron, 
in the experiment described, we might have substituted magne- 
sium or phosphorus, or indeed any one of a large number of other 
substances. In each case the result would have been similar 
to that in the case of iron and oxygen and a similar conclusion 
would have been forced upon us. In all such cases we would 
conclude that the process of burning consists in the combination 
or union of gaseous oxygen with the solid substance burned to form 
the product of the combustion. 

1 8. Burning Charcoal in Oxygen. — If we put a piece of burn- 
ing charcoal into a bottle containing oxygen we notice that it 
burns even more rapidly in oxygen than in the air. In this case 
there is but a trifling amount of ash left compared with the 
amount of charcoal burned. In order to see whether an invisible 
product may have been produced we may make the following 
experiment. If we pour a little limewater into a bottle contain- 



The Burning of Substances — Oxygen 



13 



ing oxygen and shake the limewater with the oxygen we notice 
no change. If now we pour limewater into a bottle in which 
charcoal has been burned in oxygen and again shake the con- 
tainer, the limewater becomes milky in appearance. We must con- 
clude that some invisible substance, different from oxygen, has 
been produced in the latter case. If the burning of charcoal is 
thought to be analogous in nature to the burning of iron, then 
we might expect that the product would be something composed 
of carbon and oxygen and that its weight should be equal to the 
combined weights of the carbon burned and the oxygen taken up. 
We can get some evidence that this is the case by means of the 
following experiment. 

19. Carbon Dioxide. — A small quantity of charcoal is placed 
near one end of a hard glass tube, the other end of which contains 
pieces of caustic soda (Fig. 9). If we now weigh the tube, which 



E 




Fig. 9 

may be fitted at the end containing the charcoal with a stopper 
and a small glass tube, and then cause the charcoal to burn in a 
stream of oxygen gas which we may pass through the tube, we 
shall find that there is an increase of weight, due to the fact that 
the product formed by burning the charcoal has been absorbed 
by the caustic soda in the tube. If we place some caustic soda 
in a beaker, dissolve it in water, and add some hydrochloric acid, 
we can see no marked change. If we treat the material from the 
charcoal experiment in the same way, we notice that a gas is 
given off when we pour the acid into the solution. A test of this 
gas with limewater shows that it behaves like that obtained 



14 



Introduction to General Chemistry 



when charcoal is burned directly in oxygen. The results of 
these experiments lead us to conclude that when charcoal is 
burned an invisible gas is produced, and that this gas is heavier 
than the charcoa burned; and, in fact, if charcoal had been 
burned in a closed vessel with oxygen, we should find that the 
weight of vessel and contents had not changed during the 
burning, and would be forced to conclude that the weight 'of 
the invisible product was just equal to the sum of the weights of 
the charcoal burned and the oxygen which had united with it. 
This gaseous product of the burning of charcoal was formerly 
called carbonic acid gas, but is now usually called carbon dioxide. 
20. Experiments with a Burning Candle. — We find by experi- 
ment that carbon dioxide is formed when wood, coal, illuminating 
gas, gasoline, etc., burn. We may easily show by the limewater 
test that it is also formed during the burning of a candle. We 

may also show that another well- 
known substance is produced when 
the candle burns. If we burn the 
candle under an inverted funnel 
connected by means of a glass 
tube with a U-tube which is cooled 
by immersion in a vessel of mercury 
and draw air through the funnel 
and U-tube we find that a colorless 
liquid collects in the cold U-tube 
FlG IO (Fig. 10). This liquid is water. 

The burning of the candle gives, 
then, both carbon dioxide and water. We may readily show that 
the weight of the products of a burning candle, if these are suit- 
ably collected, is greater than the weight of the candle burned. 
To do this we make use of the arrangement shown in Fig. n. 
The candle is inclosed in a glass cylinder, closed below by a cork 
having three or four holes for the admission of air. The top of 
the cylinder is filled with pieces of a solid substance (caustic 
potash) which readily absorbs both carbon dioxide and water, but 
not oxygen or nitrogen, the components of air. The apparatus 
thus arranged is suspended on one side of a balance and counter- 
poised. 



^ 




The Burning of Substances — Oxygen 



15 



The candle is now lighted and allowed to burn for ten or 
fifteen minutes, whereupon it will be found that the apparatus has 
become appreciably heavier. The increase in weight is due to the 
fact that the carbon dioxide and water formed weigh more than the 
candle burned. In fact, the excess weight is exactly the weight 
of the oxygen which has been consumed in the burning. Under 
ordinary circumstances the carbon dioxide and water escape 
our notice because both, the latter being in the form of steam, 
are invisible gases. 




Fig. 11 



21. The Law of the Conservation of Matter. — By the study 
of such facts as those discussed in the preceding paragraph and 
many others of a similar nature, Lavoisier arrived at the con- 
clusion that when a substance burns it unites with the oxygen of the 
air, and that the weight of the product is always exactly equal to the 
weight of the substance burned plus the weight of the oxygen which 
unites with the burned substance during the combustion. The 
product may be a solid, a liquid, or a gas. If it is a volatile 
liquid or a gas it usually escapes notice because it is invisible. 
Burning, therefore, consists in a union of the substance burned with 
oxygen. In this sense a substance which is burned is not destroyed; 
the material or matter composing it merely passes into another 
form, the quantity of matter in all cases being measured by its 
weight. These facts are briefly summed up in the statement that 
matter is indestructible, a statement which is frequently referred 
to as the Law of the Conservation of Matter. 



CHAPTER III 

PURE SUBSTANCES— ELEMENTS 

22. Bodies and Substances. — We use the words " substance'' 
and "body" in chemistry in very definite senses. We speak of 
things like watches or knives as "bodies." We say that the blade 
of the knife is steel, the handle is pearl. We say that a watch has 
a case of gold and a watch crystal of glass. We call steel, pearl, 
gold, and glass "substances." A substance is thus a particular 
kind of material, a body is an object which may be composed of 
one or many kinds of substances. Water, salt, and sugar are 
further examples of substances in the sense of this definition. 




Fig. 12 



23. Pure Substances. — We find that natural waters, as those 
of lakes, rivers, and springs, are not all alike. It now becomes 
important to discover the cause of the differences between 
waters 'from different sources. If we boil a quantity of lake 
water we find when the water has entirely disappeared that a 
solid residue is left. If the steam from the boiling water is con- 
densed by cooling it, as by means of a condenser (Fig. 12) through 
the outer tube of which a stream of cold water flows, we obtain 
what is called distilled water. If we now evaporate to dryness 
a quantity of this distilled water we find that no residue is left. 
If we prepare distilled water from any natural water we find that 

16 



Pure Substances — Elements 



17 



it will always evaporate completely, leaving no solid residue. 
We find further -that different kinds of natural water leave differ- 
ent proportions of solid residue upon evaporation and that the 
nature of the solid material left also differs in different cases, but 
that the distilled water in one case cannot be distinguished in 
any way from that obtained in another. We say then that 
distilled water is pure water, a pure substance, and the natural 
waters are not pure water, but that they contain dissolved foreign 
substances. If the natural water is muddy, that is, if it is not 

clear, the foreign material 
which causes it to appear 
muddy can be separated 
by filtration (Fig. 13), a 
process in which the liquid 
is allowed to seep through 
a piece of filter paper 
folded so as to fit snugly 
into a funnel. The mud 
remains on the filter paper. 
However, filtration will not 



1 



^ 




Q 



Fig. 13 



remove any of the dissolved material, but only that which is 
suspended in the water. 

24. Pure Salt Made from Rock Salt. — Common salt is found 
in nature as a mineral known as rock salt. We find that different 
samples of rock salt differ in color, taste, specific gravity, and in 
other ways. If we mix rock salt with water we find that a large 
part dissolves in the water. In general a small amount of 
material, sand, etc., will not dissolve, even though we take a large 
amount of water. If we filter the solution we separate the water 
and dissolved material from the part which has not dissolved. 
That which runs through the filter paper is called the filtrate ; it 
is the solution of the salt in water. If we boil away the water 
we find that the salt is left in the solid form and that the material 
is now free from color, that is, that it is white, and that it will 
dissolve completely in water. The salt so prepared is purer than 
the rock salt taken. Just as it is possible to prepare pure water 
from any natural water, so analogously it is possible to prepare 



Introduction to General Chemistry 



pure salt from any natural salt. Pure salt is always exactly the 
same in taste, color, specific gravity, etc., from whatever source 
it may have been obtained. The process for the purification 
of salt, described in the statement above, gives in all cases a 
much purer product than the original rock salt — pure enough 
for table use, but not a perfectly pure substance. It still con- 
tains very small amounts of some foreign substances; but even 
these can be removed by well-known methods which the student 
will learn later. A pure substance is a substance which consists 
of one sort of material. It always has definite physical proper- 
ties, from whatever source it may be obtained. 

25. Decomposition of Substances. — It was found that the red 
ash formed by heating mercury in contact with air was changed, 
upon being heated still more, into mercury and oxygen. We say 
in this case that the red ash of mercury has been decomposed into 

mercury and oxygen. We can 
accomplish the decomposition of 
many substances in an equally 
simple fashion. 

We will now consider a few 
such cases as illustrations. 

26. Decomposition of Sal 
Soda. — If we place in a test tube 
a crystal of common washing-soda, 
also known as sal soda, and heat it 
gently over a Bunsen flame, we find 
that water is produced as steam, 
and tHat it condenses in the cold 
end of the test tube. An opaque 
solid is left in place of the clear 
crystal of sal soda taken. We say 
that the sal soda has been decom- 
posed into dry soda and water. It 
would be easy to show that the weight of the water and dry soda 
formed is equal to the weight of sal soda taken. In other 
words, the sal soda has been decomposed into dry soda and 
water. 




Fig. 14 



Pure Substances — Elements 19 

27. Electrolysis of Water. — If we pass an electric current 
through some water (Fig. 14) to which we have added a few drops 
of sulfuric acid, we find that gases are produced at the platinum 
electrodes. The decomposition of a substance by an electric 
current is called electrolysis. If we collect each of these gases 
separately we find that one of them is oxygen. The other gas, 
the volume of which is double that of the oxygen, has quite differ- 
ent properties; it is called hydrogen. If we bring a lighted 
splinter into the oxygen, the splinter continues to burn with 
increased brilliancy and rapidity. If we repeat this test with 
hydrogen, we find that the hydrogen itself catches fire, just as 
illuminating gas would do, and that the splinter itself no longer 
burns in the hydrogen gas. These facts may be concisely stated 
by saying that oxygen supports combustion, while hydrogen burns 
but does not support combustion. It would be possible to show 
by experiment that the weight of the water decreases during the 
passage of the electric current through it, and that this decrease 
in weight is just equal to the combined weights of the oxygen and 
hydrogen formed. The total amount of sulfuric acid added to 
the water remains in the water at the end of the electrolysis and 
would serve to promote the decomposition of any desired amount 
of water. The complete explanation of the behavior of the 
sulfuric acid cannot be given at this point, but we know that the 
hydrogen and oxygen formed come exclusively from the water 
and not from the acid nor the platinum nor the glass of the 
vessel used. We conclude that water is decomposed by the electric 
current into hydrogen and oxygen. Therefore we may say that 
water is composed of hydrogen and oxygen or that water is a 
compound of hydrogen and oxygen. As a matter of fact, when 
hydrogen burns in air water is formed. If a cold beaker is held 
over a jet of burning hydrogen, water will be seen to condense in 
a mist on the surface of the beaker. 

28. Magnesium Burned in Steam. — That water is composed 
of oxygen and hydrogen may be shown in many other ways, one 
of which is the following. When a piece of magnesium ribbon 
burns in air the magnesium unites with the oxygen of the air 
to form a white solid which we call magnesium oxide. Now, 



20 



Introduction to General Chemistry 



magnesium will also burn in steam (Fig. 15) nearly as readily 
as it does in the air or in pure oxygen, and we find that the 
white solid which is again formed is also magnesium oxide. 
In addition, hydrogen gas is produced and may easily be col- 
lected over water. Since magnesium oxide is composed of 
magnesium and oxygen, and we obtain from magnesium and 
water magnesium oxide and hydrogen, we are again led to the 
conclusion that water is composed of hydrogen and oxygen. 

29. Steam Passed over Hot Iron. — An entirely analogous 
experiment may be carried out with iron and steam. In this 




Fig. 15 

case iron turnings or fine iron wire is strongly heated in an iron or 
glass tube (Fig. 16). When steam is passed through the tube, 
iron oxide and hydrogen are produced, a result which leads to the 
same conclusion as before regarding the composition of water. 

30. Magnesium Burned in Carbon Dioxide. — The composi- 
tion of carbon dioxide may be discovered by burning magnesium 
in this gas. We find that magnesium oxide and a product 
resembling charcoal are formed. The latter substance is carbon, 
of which charcoal is a nearly pure form. We conclude, therefore, 
that carbon dioxide is composed of carbon and oxygen or is a 
compound of carbon and oxygen. 

The facts already considered lead to the conclusion that the 
red ash obtained when mercury is heated gently in air is com- 



Pure Substances — Elements 



21 



posed of mercury and oxygen; briefly, that it is a compound of 
mercury and oxygen — a fact represented by the chemical name 
of the red ash, mercuric oxide. 

31. Elements. — The substances mercury, oxygen, hydrogen, 
and carbon have never been decomposed into simpler substances. 
We say that hydrogen and oxygen are the elements of which 
water is composed; that carbon and oxygen are the elements com- 
posing carbon dioxide. 

We may discover of what elements a substance is composed in 
two ways: either by the decomposition of the substance into the 




Fig. 16 

simpler ones which compose it — the process called analysis, or 
by causing known elements to unite in the formation of the 
original — the process called synthesis. As a result of the 
electrolysis of water we have concluded that water is composed 
of hydrogen and oxygen. This conclusion may now be tested 
by seeing whether water can be obtained from hydrogen and 
oxygen. We found that hydrogen burns readily. If we burn 
a jet of hydrogen under an inverted funnel and draw the product 
through a cooled U-tube, as in the experiment with the candle, 
we shall find that liquid water collects in the U-tube and that 
the most careful search fails to reveal any other substance as 
the product of the burning of hydrogen in air or in pure oxygen. 
Water is, therefore, a compound of the elements hydrogen and 



22 Introduction to General Chemistry 

oxygen. Since the burning of charcoal, which is a nearly pure 
form of the element carbon, gives carbon dioxide and nothing 
else, we know that carbon dioxide is a compound of the elements 
carbon and oxygen. 

32. The Burning of Copper; Copper Oxide. — If the metal 
copper, in the form of fine wire or filings, is heated in air or in 
oxygen, it is slowly changed into a black substance quite different 
in appearance from metallic copper; but during this change we 
do not observe the production of any light. By means of the 
balance we may find that the black substance formed is heavier 
than the copper taken, and we at once suspect that the copper 
has united with oxygen to form a compound. If the heating 
of the copper had been carried out in a sealed glass vessel con- 
taining oxygen, as in the earlier experiment with iron powder, 
it would have been found that gaseous oxygen had disappeared 
and that the weight of the black product was exactly equal to the 
weight of the copper taken plus the weight of the gaseous oxygen 
which had disappeared. The black substance would seem, there- 
fore, to be a compound of copper and oxygen. We know that 
when the red mercury oxide is strongly heated it is decomposed 
into mercury and oxygen. If we heat the black product from 
copper to the highest temperature we can obtain with the Bunsen 
burner, we find that it remains unaltered in weight and appear- 
ance and that no oxygen is given off. This fact might lead us 
to suspect that the black substance is not a compound of copper 
and oxygen, since its behavior is not analogous to that of mercury 
oxide. In this connection the following experiment will prove 
of interest. 

33. Hydrogen Passed over Hot Copper Oxide. — If we put two 
or three grams of the black copper product in a porcelain boat in 
a "hard" or difficultly fusible glass tube, heat the tube and con- 
tents by means of a Bunsen flame, and then pass a current of 
hydrogen through the tube, we observe that the solid glows or 
seems to burn (Fig. 17). At the same time we notice that liquid 
water condenses in the colder part of the glass tube. After a few 
minutes the glow has disappeared, even though the stream of 
hydrogen has continued. At this point the heating may be 



Pure Substances — Elements 23 

discontinued and the solid which is left in the boat allowed to 
cool in the stream of hydrogen gas. We now observe that the 
solid has the appearance and properties of metallic copper, which 
in fact it is. However, the copper is not in a single compact 
lump, for a reason which must be evident. Metallic copper can 
be melted, but the melting-point is a very much higher tempera- 
ture than that attained in the preceding experiment. Only by 
heating the copper to a point above this melting temperature 
could the material be obtained in a single lump. This could 
easily be accomplished by directing an intense blowpipe name 
upon the metal particles contained in the porcelain boat. 

We may now consider the nature of the changes which occured 
in this experiment. Since we obtained water and copper, and 



=\hk 



T7 



Fig. 17 

since we know that water is a compound of the element hydrogen 
with oxygen, we conclude that the oxygen was originally united 
with the copper and that the black substance must have been a 
compound of copper and oxygen. This substance is called 
copper oxide. We might express the result in the following 
simple fashion: Copper Oxide + Hydrogen -> Water + Copper; or 
instead of " Water" we might write " Hydrogen Oxide," the true 
chemical name for water. This statement would then show at 
a glance the nature of the chemical change which had occurred. 
34. Discovery of the Elements Composing a Substance : the 
Analysis of Malachite. — There is an almost innumerable variety 
of bodies on and in the earth, but these are composed of a very 
much smaller number of definite chemical substances. However, 
the number of definite substances is still very great, many 
thousands having been carefully described. Chemistry has for 
its object the systematic study of pure substances, their properties, 



24 Introduction to General Chemistry 

and their behavior toward one another. Happily the study of this 
immense number of substances is greatly simplified by the fact 
that they are all made up of a relatively small number of elements. 
The way in which the elements composing a substance of un- 
known composition are discovered may be illustrated by means 
of an experiment with the mineral known as malachite. Mala- 
chite is a beautiful crystalline substance often used as an orna- 
mental stone and also as one of the sources from which a familiar 
metal is obtained. If we place in a test tube, fitted with a cork 
and a bent glass tube, a few grams of malachite and heat the 
substance gently in a flame, we notice that a change in color 
from green to black occurs and at the same time that water 
condenses in the colder part of the glass tube and a gas is also 
given off. If we pass this gas into limewater we find that it 
behaves like carbon dioxide, which in fact it is. By means of 
the balance we might find that the combined weights of the 
carbon dioxide, water, and black product equal the weight of 
the original malachite. Since we know of what elements carbon 
dioxide and water are composed, it only remains to find the 
composition of the black substance in order to have a complete 
knowledge of the elements composing malachite. If this black 
substance were heated in a stream of hydrogen, it would be found 
to yield water and a red metallic-looking substance which could 
easily be recognized as copper. Therefore, the black substance 
must have been copper oxide. The results may then readily 
be interpreted. Malachite when heated is decomposed into 
carbon dioxide, water, and copper oxide. Knowing as we do the 
elements composing each of these three products, we are led to 
the conclusion that malachite is a compound of the elements 
carbon, oxygen, hydrogen, and copper. Chemists have so far been 
unable to decompose copper into anything simpler. It is, 
therefore, known as an elementary or simple substance, and we 
say that malachite is a compound of the four elements, carbon, 
hydrogen, oxygen, and copper. 

35. Some Common Elements. — The total number of known 
elements is about eighty-five, of which less than thirty are com- 
mon. In the following partial list of commoner elements, the 



Pure Substances — Elements 



25 



student will find the names of ten or twelve familiar metals. 
Carbon and sulfur, which are well known to everyone, are not 
metals; they are classed as non-metals. 

A FEW COMMON ELEMENTS 



Silver 


Copper 


Nickel 


Carbon 


Gold 


Lead 


Magnesium 


Sulfur 


Platinum 


Tin 


Zinc 


Oxygen 


Iron 


Aluminum 


Mercury 


Hydrogen 



CHAPTER IV 
THE LAW OF DEFINITE COMPOSITION 

36. The Percentage Composition of Water. — We have 
already seen that when an electric current was passed through 
water, the latter was decomposed into two gases, hydrogen and 
oxygen. It was found that the volume of the hydrogen was 
double that of the oxygen obtained in the electrolysis. This was 
not a matter of accident, for it is always found that the same 
result is obtained whenever water is electrolyzed. Since water 
is composed only of hydrogen and oxygen, we may calculate the 
percentages of hydrogen and oxygen by weight if we know the 
weight of a liter of each of these gases. Direct weighing of the 
gases has shown that 1 liter of hydrogen weighs 0.090 g. and 
1 liter of oxygen weighs 1.429 g., the gases being weighed at o° 
and 76 cm. pressure. From these figures it is easy to calculate 
that water is composed of 1 1 . 2 per cent of hydrogen and 88 . 8 
per cent of oxygen by weight. Pure water prepared from any 
source whatever always has exactly this composition. 

The percentage composition of water may also be found in 
another way. It was found in section 33 that water and copper 
are formed when hydrogen is passed over heated copper oxide. 
If this experiment be carried out with a weighed quantity of 
copper oxide, and the weight of copper which remains after the 
experiment is found, the difference in the two weights will repre- 
sent the weight of oxygen contained in the water which has been 
formed. If the weight of the water is determined, then the 
percentage of oxygen in water may readily be calculated. In 
this case we find precisely the same result as that given in the 
preceding paragraph. 

The details of the experiment are as follows. 

37. The Quantitative Synthesis of Water. — About one gram 
of pure copper oxide is placed in a weighed porcelain boat and 
heated sufficiently to drive off the moisture which it may con- 

26 



The Law of Definite Composition 



27 



tain. 1 The boat and contents are weighed as soon as cool and 
placed at once in a hard glass tube. This tube (Fig. 18) is con- 
nected at each end with U -tubes filled with calcium chloride, a 
substance that absorbs water with great readiness. One of 
these U -tubes is connected with a source of hydrogen gas and 
serves to remove all moisture (water vapor) from the hydrogen. 
The other U-tube will serve to absorb the water formed in the 
chemical reaction between the copper oxide and the hydrogen. 




WX7#'t7 



5) (^ (5} (5 




s: 



Fig. 18 

This second U-tube is accurately weighed at the beginning of 
the experiment. 

When all is ready, the stream of hydrogen is started and con- 
tinued until all air is driven from the tubes. The tube contain- 
ing the boat is now heated until the reaction begins, and kept hot 
enough beyond the boat to prevent the condensation of the 
steam formed, which is carried by the stream of hydrogen into 
the weighed U-tube. 

When all the copper oxide has been changed into copper and 
the water has all been driven over into the U-tube, the heating 
is discontinued and the copper allowed to cool in a stream of 
hydrogen. The hydrogen is then driven out by a stream of air, 
and the U-tube detached and weighed. The object in repla- 
cing the hydrogen by air is readily understood when one recalls 
that hydrogen is far lighter than air. Therefore the weight of 
the tube filled with hydrogen would be appreciably less than 
if it is filled with air. The increase in weight is the weight 
of the water formed. The boat containing the copper is also 
weighed. The loss in weight is the weight of oxygen contained 



1 Most substances, especially if porous or in the form of powder, absorb more 
or less moisture from the air. 



28 Introduction to General Chemistry 

in the water formed. The results of an actual experiment 
were as follows: 

Boat and copper oxide 9 . 523 g. 

Boat 8.451 

Copper oxide 1 . 072 g. 

Boat and copper 9 . 31 1 g. 

Boat 8.451 

Copper o . 860 g. 

Tube and water 18 . 665 g. 

Tube 18 . 426 

Water o . 239 g. 

Since 1.072 g. — 0.860 g. = 0.212 g., we conclude that o.239g. 
of water was formed from o. 212 g. of oxygen, which at the begin- 
ning was in combination with the copper in the form of copper 
oxide. Therefore water consists of 0.212 g. -5-0.239 g. =0.887 
= 88.7 per cent oxygen. The difference between the weight 
of water formed and that of the oxygen used is the weight of 
hydrogen, which is o. 239 g. — o. 212 g. =0.027 g. This is readily 
found to be n. 3 per cent of the weight of the water. Very 
carefully peformed experiments, made in this way, show that 
water contains 88 . 8 per cent by weight of oxygen and 1 1 . 2 per 
cent of hydrogen; the difference of 0.1 per cent between the 
values found in the lecture experiment quoted and those ob- 
tained in the most accurate experiments made by skilled chemists 
working with greatest care and under ideal conditions is due to 
the experimental errors in the rather crude lecture experiment. 

38. The Percentage Composition of Copper Oxide. — It is also 
easy to see that we may find the percentage composition of copper 
oxide from the data just considered. Thus 1.072 g. of copper 
oxide gave o. 860 g. of copper by loss of o. 212 g. of oxygen; from 
which we find that copper oxide is composed of 80 . 2 per cent 
copper and 19.8 per cent oxygen. The most accurate experi- 
ments made in this way give 79.9 per cent copper and 20.1 
per cent oxygen, the difference being due to experimental error 
in the lecture experiment. Pure copper oxide always has exactly 
the composition shown by these figures. 



The Law of Definite Composition 



29 



39. The Percentage Composition of Carbon Dioxide. — We 

have found that carbon in the form of charcoal burns readily in 
air or in oxygen with the formation of a colorless gas called carbon 
dioxide. The percentage composition of carbon dioxide may 
be found by burning a known weight of pure carbon in oxygen 
gas and finding the weight of carbon dioxide formed. It will be 
recalled that carbon dioxide is easily absorbed by solid caustic 
soda. It is also readily absorbed by a solution of caustic potash 
in water, while neither oxygen nor air is absorbed by such a solu- 
tion. If the gases formed by the burning of carbon in a stream 
of oxygen are passed through a suitable bulb containing caustic 
potash solution, all of the carbon dioxide will be retained by the 
solution and the oxygen will pass through unabsorbed. The 
increase in weight of the bulb will represent the weight of the 
carbon dioxide formed by the burning of the carbon. * 




Fig. 19 

The arrangement of the apparatus is shown in Fig. 19. 
About o . 2 g. of pure carbon, made from sugar, is contained in a 
porcelain boat which is placed in a hard glass tube connected at 
one end with a supply of pure oxygen and at the other with a 
calcium chloride tube and a weighed potash bulb, which contains 
a 30 per cent solution of caustic potash. The middle part of the 
tube should contain a column of copper oxide, to insure the 
complete conversion of the carbon into carbon dioxide. The 
calcium chloride tube serves to catch any moisture present. 
The carbon is ignited by heating the tube with a gas burner; 
after the carbon has completely burned and all of the carbon 
dioxide formed has been driven over into the potash bulb by 
the stream of oxygen, a slow stream of air is blown or drawn 
through the apparatus to replace the oxygen by air. The 
potash bulb is then detached and weighed. In an actual lecture 
experiment o. 194 g. of carbon yielded o. 701 g. of carbon dioxide; 



30 Introduction to General Chemistry 

from which we find that this gas contains 27.6 per cent of carbon 
and 72.4 per cent of oxygen. The most accurate experiments 
of skilled chemists show the correct percentages to be 27.3 per 
cent carbon and 72. 7 per cent oxygen. 

40. The Action of Sodium on Water: Caustic Soda. — It is a 
matter of importance to know the exact percentage composition 
of pure substances and a great variety of methods must be 
employed in the making of such determinations. It often 
happens that the method which would seem to be most direct 
and desirable is not practicable because the violence of the inter- 
action of the elements which we bring together would cause loss 
of some of the material taken. This may be illustrated by an 
experiment with the element sodium. If we throw a piece of 
this metal upon water, we observe that the action is a violent 
one which 'ordinarily ends in an explosion that throws part of 
the substance out of the beaker in which it was contained. We 
may carry out the same reaction without loss of material and 
obtain precisely the same product if the piece of sodium is 
exposed to water vapor instead of being thrown upon liquid 
water. In this case the reaction requires much more time, but 
it proceeds quietly and without loss of material. The white 
solid so obtained is caustic soda. 

41. The Action of Hydrochloric Acid on Caustic Soda: Com- 
mon Salt. — If we add to a solution of caustic soda contained in a 
beaker a sufficient amount of pure hydrochloric acid and evap- 
orate the resulting solution to dryness, we find that the product 
is one with which we are well acquainted. It is nothing more 
nor less than common salt, and if the materials used are all pure 
the product will be chemically pure salt. We discover in this 
way that the metal sodium is one of the constituents of common 
salt. In fact, metallic sodium may be obtained by the elec- 
trolysis of molten salt, although this is not the most satisfactory 
method of making this metal. The percentage of sodium in salt 
may readily be found if the weights of sodium taken and of salt 
obtained are determined. 

42. The Percentage of Sodium in Common Salt. — In an 
actual experiment o . 483 g. of metallic sodium was weighed in a 



The Law of Definite Composition 31 

stoppered test tube (to prevent action of the moisture of the air) . 
The sodium was placed on a strip of silver foil which rested on the 
edges of a small porcelain dish containing about 10 c.c. of water, 
and covered with a beaker. In the course of a few hours the 
sodium had reacted completely with the water vapor to form a 
solution of caustic soda which dripped into the dish. A little 
of the solution adhering to the foil was rinsed into the dish with 
a little water. Sufficient pure hydrochloric acid was then added 
and the solution evaporated by steam heat in the manner shown 
in Fig. 20. The beaker contained ordinary water. By this 
mode of evaporation of the solution in the dish we avoid loss by 
spattering that would occur if we should 
boil the solution by heating the dish 
directly with the flame. When the salt 
appeared to be dry, the dish was heated 
very cautiously with the direct flame, to 
drive off the small amount of remaining 
water. When cold, the dish and contents 
were weighed. It was found in this way — 



JM 



that 0.483 g. of sodium gave 1.217 g. of p IG# 20 

common salt, which indicated that salt 
contains 39.7 per cent of sodium. The correct result is 39.4 
per cent. 

43. The Electrolysis of Hydrochloric Acid: Chlorine. — It is, 
of course, obvious that the sodium in common salt must be com- 
bined with one or more elements and the student will readily 
guess that a clue to the other constituents of common salt may 
be gained by a knowledge of the constituents of hydrochloric 
acid. If we pass an electric current through a concentrated solu- 
tion of hydrochloric acid contained in the apparatus shown in 
Fig. 21, we find that two gaseous products are obtained, the 
volumes of which are practically equal. One of these is colorless. 
It is lighter than air and burns with a hot but non-luminous 
flame and in so doing yields water; these properties show the 
colorless gas to be hydrogen. The other gas is pale yellow in 
color; it is heavier than air, one liter weighing 3. 22 g., and has 
an exceedingly disagreeable, irritating odor. This gas is known 



32 



Introduction to General Chemistry 



as chlorine. Inasmuch as chlorine has never been separated 
into simpler substances, we conclude that it is an element. 

44. The Union of Hydrogen and Chlorine: Hydrogen 
Chloride Gas. — Since the hydrochloric acid which was elec- 
trolyzed contained water, we should not be warranted in con- 
cluding that hydrogen is a constituent of hydrochloric acid; 

for, as we know, hydrogen is also one of 
the constituents of water. If we bring 
together equal volumes of the gases 
hydrogen and chlorine and allow them 
to mix, and if we allow the vessel to 
stand in diffused light for a day or 
two, we notice that the yellow color 
of the chlorine has disappeared. We 
find that a colorless gas remains which 
dissolves with the greatest ease in 
water, and that neither hydrogen nor 
chlorine is left. Since water which has 
dissolved the gas has all of the proper- 
ties of a solution of pure hydrochloric 
acid, we interpret the results as show- 
ing that equal volumes of hydrogen 
and chlorine gases combine to form 
a new gas which we call hydrogen 
chloride gas, and that the latter when 
dissolved in water constitutes hydro- 
chloric acid. Hydrogen chloride gas may be distinguished from 
the other gases which we have met in several ways, notably by 
its marked, choking odor, by the fact that it fumes or gives a 
white cloud in moist air, and it dissolves with great ease in 
water, as well as in several other ways. 

45. Salt a Compound of Sodium and Chlorine. — The fact that 
hydrochloric acid is known to be a compound of chlorine suggests 
that common salt may also contain this element. This is in 
fact the case. It can readily be shown by experiment that 
common salt results from the union of chlorine gas with metallic 
sodium. Inasmuch as nothing else is needed and no other 




Fig. 21. 

apparatus. 



Brownlee's 



The Law of Definite Composition 33 

product than salt is formed, we must conclude that salt is a 
compound of the elements sodium and chlorine. This fact is 
indicated by the chemical name of common salt, sodium 
chloride. Since salt contains 39.4 per cent of sodium, the 
percentage of chlorine must be 60.6. 

46. The Law of Definite Composition. — The preceding para- 
graphs of this chapter are intended to illustrate how we may 
arrive at a knowledge of the nature and percentage by weight 
of each element entering into the composition of a pure substance. 
It is possible, by well-known methods, to do this for all pure 
substances. As a result of countless thousands of such quantita- 
tive experiments made by chemists, the conclusion has been 
reached that the percentage composition of every pure substance 
is perfectly definite for that substance and is found to be the same 
by whatever method we may make the determination. This is 
one of the most important laws of chemistry. It is usually 
spoken of as the Law of Definite Composition or of Definite 
Proportions. This explains why a pure substance always has 
definite properties, from whatever source it may be obtained. 

47. Hydrogen and Its Gaseous Compounds. — We have 
already become acquainted with hydrogen and one of its gaseous 
compounds, hydrogen chloride, a water solution of which is 
known as hydrochloric acid. Hydrogen forms many compounds 
which are gaseous at ordinary temperatures. We shall now take 
up a study of some of these, with the object in view, first, of dis- 
covering the nature of the other element combined with the 
hydrogen; secondly, of discovering the percentage composition; 
and, finally, of disclosing a very remarkable relation between the 
weights of hydrogen contained in equal volumes of these gases. 

48. Hydrogen Chloride. — We have found that equal volumes 
of hydrogen and chlorine combined to form hydrogen chloride 
gas. Since we know that 1 liter of hydrogen weighs 0.090 g. 
and that 1 liter of chlorine weighs 3. 220 g., we find by calcula- 
tion that hydrogen chloride contains 2.76 per cent by weight of 
hydrogen. By direct weighing of pure hydrogen chloride gas 
it is found that 1 liter weighs 1.642 g. Since 2.76 per cent of 
1.642 g. is 0.045 g-> ^ follows that 1 liter of hydrogen chloride 



34 Introduction to General Chemistry 

gas contains 0.045 g. of combined hydrogen. It has already 
been stated that 1 liter of hydrogen gas weighs 0.090 g., which 
weight we see is exactly double the weight of hydrogen in 1 liter of 
hydrogen chloride gas. 

49. Acetylene: a Compound of Carbon and Hydrogen.— 
Let us next consider the gas acetylene which is extensively used 
for illumination. This gas is obtained by allowing water to 
drop on calcium carbide. We find that it is a colorless gas with 
a peculiar odor. Everyone knows that it burns in air, giving an 
exceedingly bright flame. If we collect and test the products 
coming from the acetylene flame we find carbon dioxide and 
water. We find the same products and no others when acetylene 
is burned in pure oxygen gas, and therefore conclude that carbon 




Fig. 22 

and hydrogen are constituents of acetylene; but the experiment 
obviously does not decide whether oxygen is or is not also a 
constituent of acetylene. This question could be decided if we 
knew the percentages of carbon and hydrogen in the gas. 

50. The Analysis of Acetylene. — We may find the per- 
centages of carbon and hydrogen by means of the following 
experiment. A tube of hard glass a centimeter or more in 
diameter and 30 cm. long (Fig. 22) is partly filled with pure dry 
copper oxide. The tube is then heated red hot and a measured 
volume of acetylene at a known temperature and pressure is 
caused to pass through the tube and over the heated copper 
oxide. It is found that carbon dioxide and water are formed 
and that part of the copper oxide is changed into metallic copper. 
A U-tube filled with calcium chloride, for the absorption of the 
water formed, is attached to the exit of the hard glass tube. 
Beyond this, attached by rubber tubing, we have a bulb contain- 
ing caustic potash solution to absorb the carbon dioxide. After 



The Law of Definite Composition 35 

all of the acetylene has been driven over into the combustion 
tube holding the copper oxide, by allowing mercury from the 
attached reservoir slowly to displace the acetylene, a slow stream 
of pure dry oxygen is passed into the combustion tube to insure 
the complete burning of the carbon of the acetylene. Finally, 
the oxygen is displaced by a stream of air. 

The increase in weight of the calcium chloride tube represents 
the weight of water formed. Similarly the increase in weight 
of the caustic potash bulb represents the weight of carbon dioxide 
obtained. Now we know that water contains 1 1 . 2 per cent of 
hydrogen and that carbon dioxide contains 27.3 per cent of 
carbon. We may then calculate the weights of hydrogen and 
carbon corresponding to the weights of water and carbon dioxide 
obtained. If we know that 1 liter of acetylene under standard 
conditions, that is, at o° and 76 cm. P, weighs 1 . 190 g., we have 
all the data needed to enable us to calculate the percentages of 
hydrogen and carbon in acetylene. In an actual lecture experi- 
ment 200 c.c. of pure dry acetylene at 18 and 75.4 cm. gave 
o.i5og. of water and 0.751 g. of carbon dioxide. From the 
data above we find that the weight of the acetylene taken was 
o. 222 g., and that the weights of hydrogen and carbon contained 
in the water and carbon dioxide respectively were 0.0168 g. and 
0.205 £•> respectively. Therefore acetylene contains (according 
to this analysis) 7.5 per cent of hydrogen and 92.3 per cent of 
carbon. The correct percentages are 7.7 and 92.3 respectively; 
and since the sum of these percentages is 100, we know that 
hydrogen and carbon are the- only elements contained in acetylene. 
We may also calculate from the same data the weight of com- 
bined hydrogen in one liter of acetylene under standard condi- 
tions. We find in this way o . 090 g. of hydrogen. 

51. Ammonia. — Let us next take up the study of ammonia. 
Common household ammonia, which is familiar to everyone, is a 
solution in water of the substance, ammonia, which is a gas at 
ordinary temperature and pressure. If we warm such a solution 
of ammonia, a gas having an intense odor is given off. When 
this gas, ammonia, is strongly compressed, it condenses to a 
colorless liquid which we speak of as liquid ammonia. This is 



36 Introduction to General Chemistry 

a commercial article which is shipped in heavy steel cylinders six 
feet long and a foot in diameter. The liquid ammonia exists 
under considerable pressure in such cylinders. If the valve of 
the cylinder is opened gaseous ammonia escapes. We may use 
a small cylinder of liquid ammonia as a convenient source of 
ammonia gas. 

If we fill a glass cylinder with mercury, invert it in a dish of 
mercury, and allow ammonia gas to escape under the mouth of 
the cylinder, the mercury is displaced by the ammonia gas. We 
notice that the gas is invisible, like air. It is to be distinguished 
from air, however, by its intense odor, as well as in other ways. 
If we dip the mouth of the cylinder, which has been closed by a 
glass plate, into a vessel of water, we find that the water rushes 
into the cylinder almost as readily as if the space were a vacuum. 
An examination of the water now shows that it has new prop- 
erties. The water now has the odor of ammonia, it has a 
peculiar disagreeable taste, and changes the color of immersed 
red litmus paper blue. If we bring a burning candle into a 
cylinder of ammonia the flame of the candle is extinguished but 
the ammonia does not take fire. These properties distinguish 
ammonia from oxygen, hydrogen, and acetylene. 

52. Ammonia a Compound of Nitrogen and Hydrogen. — We 
may now inquire, What is the chemical composition of ammonia? 
Is it an elementary substance or a compound, and, if a compound, 
of what elements is it composed? If ammonia gas is passed 
through a heated glass tube containing copper oxide we observe 
that a colorless liquid condenses in the cold part of the tube. 
This liquid proves to be water. We find also that a colorless, 
odorless gas is formed. If we pass this gas into limewater we 
observe no result and conclude, therefore, that this gas is not 
carbon dioxide. We find that the gas is not appreciably soluble 
in water, so that it cannot be unchanged ammonia gas. If we 
test the gas with a burning candle we find that it neither burns 
nor supports combustion. The student will doubtless recall 
(10) that this gas has just those properties which the portion 
of the air left after the removal of oxygen by mercury or phos- 
phorus possesses. It would seem, therefore, to be nitrogen. 



The Law of Definite Composition 37 

The identity of the gas with nitrogen is confirmed by a deter- 
mination of the density; whereupon it is found that a liter 
weighs i.25ig. Since water and copper were formed from 
ammonia and copper oxide, we conclude that ammonia has 
furnished the hydrogen which united with the oxygen supplied 
by the copper oxide to form the water obtained in the preceding 
experiment. Ammonia must be a compound containing nitrogen 
and hydrogen. It has been shown in many ways by experiments, 
which we need not consider at present, that nitrogen and hydro- 
gen are the only constituents of ammonia. 

53. The Percentage Composition of Ammonia. — The per- 
centage of hydrogen in ammonia may be found by carrying 
out the experiment above described with a known volume of 
ammonia measured at a known temperature and pressure. If 
we cause the ammonia to pass through the heated copper oxide 
tube, driving out water vapor completely by means of air after 
all of the ammonia has passed into the tube, and if the products 
are caused to pass through a calcium chloride tube connected 
to the copper oxide tube as in the determination of the composi- 
tion of acetylene, the increase in weight of the calcium chloride 
tube gives us the weight of water formed from the hydrogen of 
the ammonia used. Knowing as we do the percentage of hydro- 
gen in water, if we know the weight of a liter of ammonia gas 
(o.772g.) we may calculate the percentage of hydrogen in 
ammonia and also the weight of combined hydrogen in 1 liter of 
ammonia gas measured under standard conditions. We find 
this latter weight to be o. 135 g. 

54. Methane, Another Compound of Carbon and Hydro- 
gen. — The chief component of natural gas is a substance called 
methane. This same gas methane often escapes in bubbles when 
the decaying vegetable matter in marshes is disturbed. For this 
reason methane is also known as marsh gas. We may prepare 
methane artificially in the laboratory by methods which we need 
not now discuss. It may be collected over water, as its solu- 
bility in water is slight. We note that it is a colorless gas, that 
it is lighter than air, since the gas will escape rapidly from an 
open cylinder when the mouth of the cylinder is turned upward, 



38 Introduction to General Chemistry 

but will not escape if the mouth is downward. One liter of 
methane weighs 0.721 g., which is but little more than half of 
the weight of the same volume of air. If we bring a lighted 
candle into a cylinder of methane we find that the gas burns 
with a slightly luminous flame but that the candle flame is 
extinguished. 

55. The Quantitative Analysis of Methane. — If we examine 
the products of combustion from a methane flame we find water 
and carbon dioxide, from which we know that methane is a com- 
pound of carbon and hydrogen with or without other elements. 
We may determine the quantitative composition of methane by 
precisely the same method as that used for the quantitative 
analysis of acetylene, whereupon we find that methane contains 
75.0 per cent of carbon 'and 25.0 per cent of hydrogen by weight. 
Since the sum of these percentages is 100 we know that methane 
must contain only the elements carbon and hydrogen. From the 
data obtained in the analysis of methane we may also calculate 
that 1 liter of methane under standard conditions contains 
o. 180 g. of combined hydrogen. 

56. The Weight of Hydrogen in One Liter of Gaseous Hydro- 
gen Compounds. — By a study of the composition of the four 
gases, hydrogen chloride, acetylene, ammonia, and methane, as 
well as of hydrogen itself, we have found the weight of hydrogen 
in 1 liter of each. These results may now be tabulated as in 
Table II. An inspection of the results given in the table reveals 

TABLE II 

Hydrogen chloride o . 045 g. 

Hydrogen o . 090 

Acetylene o . 090 

Ammonia 0.135 

Methane o. 180 

a remarkable fact. The weight of hydrogen in 1 liter of hydrogen 
chloride is less than that in any other case. The weight per liter 
of hydrogen gas itself is double the weight of Hydrogen in 1 liter of 
hydrogen chloride. Likewise the weight of hydrogen in 1 liter of 
acetylene is exactly equal to the weight of a liter of free hydrogen and 



The Law of Definite Composition 39 

also double the weight of hydrogen in 1 liter of hydrogen chloride. 
The weight of hydrogen in 1 liter of ammonia is three times that in 
1 liter of hydrogen chloride, while in the case of methane the weight 
of hydrogen per liter is four times the weight of this element in the 
same volume of hydrogen chloride. 

If we consider the weight of hydrogen in a liter of hydrogen 
chloride as unity, we find that the weights in the same volumes of the 
other gases are expressed by the numbers 2, 3, or 4. It is obvious 
that the relations we discussed would also hold equally well if 
we dealt with weights of hydrogen contained in any other fixed 
volume, as a cubic foot or a cubic meter. We could express the 
facts by saying that the weight of hydrogen contained in a fixed 
volume of any of these gases is in each case a multiple of the mini- 
mum weight, which is found in the case of hydrogen chloride gas. 
Since 1 liter of hydrogen chloride gas contains 0.045 g. of hydro- 
gen, 1 g. of combined hydrogen would be contained in 22.4 
liters 1 of hydrogen chloride. In the same volume of the other gases 
the weights of hydrogen would be 2 g., 3 g., or 4 g. 

1 In reality i-f- 0.045 gives 22.2 instead of the correct value 22.4 liters. The 
discrepancy is caused by the fact that the members used are only approximate. 
This subject is discussed further in section 222. 



CHAPTER V 
SYMBOLS AND CHEMICAL FORMULAE 

57. Gaseous Carbon Compounds. — We may now inquire 
whether the remarkable relations between the weights of hydro- 
gen in equal volumes of compounds of hydrogen hold good in the 
case of compounds of other elements. We have already studied 
three gaseous compounds of carbon: carbon dioxide, acetylene, 
and methane, and have seen how the percentage composi- 
tion of each is determined. Before discussing the results so 
obtained, let us consider two new gaseous compounds of carbon : 
propane and trimethylamine. 

58. Propane: a Compound of Carbon and Hydrogen. — 
Propane is found in small amounts in the natural gas of some 
wells and also dissolved, in small quantities, in crude petroleum. 
It may also be obtained artificially by methods well known to the 
chemist, the nature of which we need not now consider. We 
observe that propane is a colorless, odorless gas which is some- 
what heavier than air, 1 liter under standard conditions weighing 
1.97 g. We find that propane resembles methane in its chemi- 
cal behavior, since it extinguishes a burning candle but takes 
fire itself at the same time, burning with a slightly luminous 
flame and yielding carbon dioxide and water as the only products 
of combustion. The analysis of propane may be carried out in 
precisely the same manner as our analysis of methane and acety- 
lene. We find in this way that propane contains 81.8 per cent 
of carbon and 18.2 per cent of hydrogen. Since the sum of these 
percentages is 100, it follows that carbon and hydrogen are the 
only constituents of propane. 

59. Trimethylamine : a Compound of Carbon, Hydrogen, and 
Nitrogen. — Trimethylamine is a colorless gas about twice as 
heavy as air, 1 liter weighing 2.65 g. Its odor is very powerful 
and somewhat disagreeable, but if inhaled in small quantities the 
gas is not poisonous nor irritating, as is, for example, chlorine gas. 
The odor is that of decaying fish. In fact, the gas can be 

40 



Symbols and Chemical Formulae 



4i 



obtained from products separated from herring brine. We find 
that the gas is very easily soluble in water and that the solution 
turns red litmus paper blue, just as ammonia does; but the gas 
may be distinguished from ammonia by the fact that it will burn, 
whereas ammonia will not. It is easy to discover that water and 
carbon dioxide are formed when trimethylamine is burned in air 
or in oxygen. If we pass trimethylamine through a tube con- 
taining heated copper oxide we obtain, in addition to water and 
carbon dioxide, a colorless, odorless, incombustible gas which can 
easily be identified as nitrogen. These facts show that tri- 
methylamine contains the elements carbon, hydrogen, and nitrogen. 
We could determine the percentages of carbon and hydrogen by 
finding the weights of carbon dioxide and water formed by the 
action of the gas on hot copper oxide, as in analyses previously 
made. We might also find the percentage of nitrogen by finding 
the volume of nitrogen which we could obtain from a known 
volume of the gas. The percentages of carbon, hydrogen, and 
nitrogen would be found to be 61 . o, 15.3, and 23 . 7 respectively. 
60. The Weights of Carbon in 1 Liter and in 22.4 Liters of 
Gaseous Carbon Compounds. — Let us now consider the facts 
presented in Table III. The weight of 1 liter and the percentage 

TABLE III 





Weight of 
1 Liter 


Percentage of 
Carbon 


Weight of 

Carbon in 

1 Liter 


Weight of 
Carbon in 
22.4 Liters 


Methane 

Carbon dioxide 

Acetylene 

Propane 

Trimethylamine 


0.72 

1-97 
1. 19 
I.97 

2.65 


75-0 
27-3 
9 2 -3 
81.8 
61 .0 


0-54 
0.54 
1.08 
1.62 
I.62 


12 
12 
24 
36 
36 



of carbon in each of the five gaseous compounds of carbon we 
have studied are given in the first and second columns of figures. 
The product of the weight of 1 liter of a gas by the percentage of 
carbon it contains gives the weight of combined carbon in 1 liter. 
These products are given in the third column. The weights of 
carbon in 22.4 liters, as given in the last column, are found by 
multiplying the corresponding weights in the third column by 
22.4. 



42 Introduction to General Chemistry 

We see by a glance at the last column of the table that 22.4 
liters of carbon dioxide and methane contain 12 g. of combined 
carbon, that the same volume of acetylene contains 24 g. of 
carbon, while the weight of combined carbon in 22.4 liters of 
propane and trimethylamine is 36 g., and therefore that the 
weight of carbon in 22 . 4 liters of any of these gases is either one, 
two, or three times 12 g. In the case of gaseous hydrogen com- 
pounds, we found that the weight of hydrogen was either one, 
two, three, or four times 1 g., which was the minimum weight 
of this element found in any case. We thus find that in 22.4 
liters of various pure gases the minimum weight of hydrogen is 1 g. 
and the minimum weight of carbon 12 g., and, further, that if a 
greater weight of either of these elements is contained in this volume 
of any pure gas, the weight is a multiple of the minimum weight 
by a small whole number. 

Let us now consider the weights of carbon and hydrogen 
contained in 22.4 liters of the three gaseous compounds which 
contain only carbon and hydrogen, namely, methane, acetylene, 
and propane. In 22.4 liters of methane we find 12 g. of carbon 
combined with 4 g. of hydrogen. In the same volume of acety- 
lene, 24 g. of carbon combined with 2 g. of hydrogen, and in the 
case of propane 36 g. of carbon combined with 8 g. of hydrogen. 
Without considering at present the theoretical significance of 
the remarkable facts which these figures show, we may consider 
a practical application of the facts which will enable us to express 
the composition of these gases in a simple fashion. 

The student must realize that since we have three compounds 
all consisting of carbon and hydrogen and having different prop- 
erties, the difference in percentage composition must be an 
important factor in determining the properties of the substance. 
He will also understand that a knowledge of the percentage com- 
position is a matter of prime importance for the chemist, and that 
any scheme by means of which a knowledge of the composition 
by weight could be easily memorized would be important. 

61. Symbols. — Suppose we represent 1 g. of hydrogen by a 
sign or symbol and choose the letter H for this purpose. We 
could, then, represent by H taken four times the weight of hydro- 



Symbols and Chemical Formulae 43 

gen contained in 22.4 liters of methane; by H taken twice, or 
2H, the amount of hydrogen in 22.4 liters of acetylene; and 
similarly by 8H, the amount of hydrogen in 22 . 4 liters of propane. 
Suppose that, on the other hand, we represent 1 2 g. of carbon by 
the sign or symbol C, then C, 2C, and 3C will represent the 
weights of carbon in 22.4 liters of methane, acetylene, and 
propane respectively. The weights of carbon and hydrogen in 
22.4 liters of methane may then be represented by writing iC 
together with 4H. As a matter of convenience the multiples 1 
for the C and 4 for the H, are written as subscripts; so that 
instead of iC and 4H we write dH 4 . In practice no subscript 
is used when the multiple is 1 . The composition of methane is 
represented simply by CH 4 . 

62. Chemical Formulae.— In a similar way we may represent 
the weights of carbon and hydrogen in 2 2 . 4 liters of acetylene by 
C 2 H 2 while the composition of the same volume of propane may 
be represented by C 3 H 8 . We call H the symbol for hydrogen, and 
for the present we may consider that H or iH represents 1 g. 
of hydrogen and similarly that C, the symbol for carbon, repre- 
sents 12 g. of that element. We call the expressions CH 4 , C 2 H 2 , 
and C 3 H 8 the formulae of methane, acetylene, and propane 
respectively. We shall now proceed to show how this system 
may be extended to all gaseous compounds of any element what- 
ever. 

Chemists are familiar with a large number of gases in addi- 
tion to those which we have already studied. Some of these are 
of much practical importance while others are chiefly of interest 
to the chemist for scientific reasons. In every case it is a simple 
matter to determine the weight of 1 liter of the gas under stand- 
ard conditions, the method of making the determination being 
essentially the same in all cases. Furthermore, by methods 
which are well known to chemists we may determine what ele- 
ments compose any gas, and by means of a quantitative analysis 
we may determine the percentage of each element in the gas. 
If we calculate in the case of each gas the weight of each element 
contained in 22.4 liters of the gas, we obtain results like those 
shown in Table IV. 



44 



Introduction to General Chemistry 



63. The Minimum Weights of Oxygen, Nitrogen, and 
Chlorine. — An inspection of the results given in Table IV shows 
that the same regularity in the weights of hydrogen and carbon 
holds in all cases, as we have observed it to hold in the few cases 
discussed in the preceding paragraphs. We notice also that the 
minimum weight of oxygen in 22.4 liters of any of its gaseous 

TABLE IV 

Weights of Constituents in 22.4 Liters of Gases 



Substance 



Oxygen 



Hydrogen 



Carbon Nitrogen Chlorine Formula 



Oxygen 

Carbon monoxide . . 
Carbon dioxide .... 

Nitrous oxide 

Nitric oxide 

Nitrosyl chloride . . . 
Hypochlorous oxide. 
Chlorine dioxide . . . 

Phosgene 

Methyl ether 

Hydrogen 

Hydrogen chloride. . 

Prussic acid 

Ammonia 

Methane 

Acetylene 

Ethylene 

Ethane 

Propylene 

Propane 

Methyl chloride 

Ethyl chloride 

Methylamine 

Nitrogen 

Cyanogen 

Cyanogen chloride. . 

Chlorine 

Trimethylamine. . . . 



2X16 
1X16 
2X16 
1X16 
1X16 
1X16 
1X16 
2X16 
1X16 
1X16 



1X12 
1X12 



2X14 
1X14 
1X14 



6X1 
2X1 
1X1 
1X1 
3X1 
4X1 
2X1 
4X1 
6X1 
6X1 
8X1 
3Xi 
5Xi 
5Xi 



1X12 
2X12 



1X12 



1X14 
1X14 



1X12 
2X12 
2X12 
2X12 
3X12 
3X12 
1X12 
2X12 
1X12 



2X12 
1X12 



1X14 
2X14 
2X14 
1X14 



9X1 



3X12 



1X14 



1X35-5 
2X35-5 
1X35-5 
2X35-5 



iX35-5 



1X35-5 
iX35-5 



iX35-5 
2X35-5 



2 
CO 

co 2 

N 2 

NO 

NOC1 

C1 2 

C10 2 

COCl 2 

C 2 H 6 

H 2 

HC1 

HNC 

NH 3 

CH 4 

C 2 H 2 

C 2 H 4 

C 2 H 6 

C 3 H 6 

C 3 H 8 

CH3CI 

C 2 H S C1 

CH S N 

N 2 

C 2 N 2 

C1NC 

Cl 2 - 

C 3 H 9 N 



compounds is 16 g., and that this weight is found in many 
cases, while in others the weight is twice 16. In the case of 
the compounds of nitrogen we note that the minimum weight 
is 14 g. and that in other cases the weight is double this mini- 
mum weight. In the case of chlorine compounds the minimum 
weight of chlorine is 35.5 g., while those compounds with 
a larger proportion of chlorine contain double the minimum 
weight. 



Symbols and Chemical Formulae 45 

64. The Law of Minimum and Multiple Weights. — Entirely 
analogous regularities will be found if we consider the data 
obtained from a study of the gaseous compounds of any other 
elements. For each element we find a minimum weight in the 
volume of 22 .4 liters of any of its gaseous compounds under standard 
conditions and also find that the weight if greater than the minimum 
would be 2, j, or 4, or some small multiple of this minimum. This 
last statement may be called the Law of Minimum and Multiple 
Weights. 

65. The Chemical Unit Volume: 22.4 Liters. — The volume 
22.4 liters thus becomes a kind of unit volume for the chemist, 
this particular volume having been chosen because it contains 
1 g. of hydrogen in the case of those hydrogen compounds 
which contain the minimum weight of this element. In this 
volume no other element has a minimum weight as small as that 
of hydrogen. 

66. Symbols Represent Minimum Weights. — In the same 
manner as that suggested in a preceding paragraph for hydrogen 
and carbon, we may represent the minimum weight of each of 
the other elements by a symbol. Table V shows the minimum 
weights of the five elements we have been considering, together 
with the corresponding symbols. 

TABLE V 

Minimum Weights in 22.4 Liters, and Symbols 



Hydrogen , 
Carbon. . . 
Nitrogen . 
Oxygen. . 
Chlorine. . 




H 

C 
N 
O 
CI 



67. Making Formulae. — We see from Table IV that 22.4 
liters of carbon dioxide contain 12 g. of carbon combined with 
2X16 g. of oxygen. We may, therefore, represent the composi- 
tion of the quantity of carbon dioxide in 22.4 liters by the 
formula C0 2 . In an analogous fashion we may obtain as the 
formula representing the composition of 22.4 liters of ammonia, 
NH 3 , and as the formula for hydrogen chloride, HC1. By making 



46 Introduction to General Chemistry 

use of this system the student will now have no difficulty in 
writing down at once the formula of each of the gases from the 
data contained in Table IV. He will also readily see that it is a 
much less difficult task to learn the formulae of such gases than 
to learn their percentage composition; that is to say, it is an easier 
tax upon the mind to remember the formula HC1 than to remem- 
ber that hydrogen chloride contains 2.76 per cent of hydrogen 
and 97 . 24 per cent of chlorine. 

68. The Practical Use of Formulae. — A review of the methods 
employed in arriving at the results represented by the formula 
of any substance shows that in each case we have made use of the 
knowledge, first, of the weight of 1 liter of the gas, and, secondly, 
of the percentage of each of its elementary constituents. Con- 
versely, if we know the weights which the symbols of the elements 
represent, and know the formula of a gas, we may by working 
backward find its percentage composition. For example, sup- 
pose that we remember that the formula of methane is CH 4 
and know that H stands for 1 g. of hydrogen and C for 12 g. 
of carbon. Then 22.4 liters of methane contain 4 g. hydrogen 
combined with 12 g. of carbon. The proportion of hydrogen is, 
therefore, 4/16, or 25 per cent, and of carbon 12/16, or 75 per cent, 
the weight of 22 .4 liters being 16 g., and 1 liter weighs 16/22.4= 
0.72 g. In calculating in this way the density and percentage 
composition of methane we are merely reproducing the results 
which originally were obtained by experiment. In order to find 
the formula of any gas, we must know its density and the per- 
centage of each elementary constituent. We find by actual 
experience that we can represent by a formula, usually of a very 
simple character, the composition of 22.4 liters of any gaseous 
substance. 

69. Formulae of Liquids and Solids. — The system which we 
have just considered is capable of extension to liquid and solid 
substances, in which case, however, the formula may have a 
slightly less definite meaning. We may illustrate this by con- 
sidering the cases of water and mercury oxide. We have found 
that water is composed of 1 1 . 2 per cent of hydrogen and 88 . 8 
per cent of oxygen, from which we observe that the weight of 



Symbols and Chemical Formulae 47 

the oxygen is 8 times the weight of the hydrogen with which it is 
united. This ratio of hydrogen and oxygen might be represented 
by H 2 0, since this formula would mean that 2 g. of hydrogen are 
united with 16 g. of oxygen, which weights of hydrogen and 
oxygen are in the ratio of 1 to 8, but the formulae H 4 2 and 
H 6 3 would also represent equally well the proportion of hydro- 
gen and oxygen actually found in water. 

70. The Formula of Water. — We may be led to choose a con- 
sistent formula for water by the consideration of the density of 
water vapor or steam; but in this case the density determination 
must be made at a temperature above the boiling-point of water, 
if we work at atmospheric pressure. Since the effect of changes 
of pressure and temperature upon the volume of a given quantity 
of steam are the same as upon an equal volume of any gas which 
would not liquify if cooled to o° at 76 cm. pressure, we might 
calculate by the laws of Boyle and Charles what the volume of 
the known weight of steam measured at a high temperature and 
known pressure would be if the steam were under standard con- 
ditions, that is, at o° and 76 cm. pressure. It has been found in 
this way that 1 liter of water vapor if it did not condense to a 
liquid would weigh o . 806 g. under standard conditions, which 
corresponds to a weight of 18 g. for 22.4 liters. Now, n . 2 per 
cent of 18 g. is 2 g. and 88.8 per cent of 18 g. is 16 g. From 
these results we conclude that if water vapor could exist under 
standard conditions as a gas that 22.4 liters would contain 
2 g. of combined hydrogen and 16 g. of combined oxygen, which 
amounts would be exactly represented by the formula H 2 0. 

71. Formulae of Volatile Liquids and Solids. — In a perfectly 
analogous fashion we could find the formula for any other volatile 
substance, the density of whose vapor we could measure experi- 
mentally. Such a procedure would enable us to represent by a 
formula the composition of a great number of volatile chemical 
substances which are not gaseous, but are liquid or solid under 
ordinary conditions of temperature and pressure. 

72. Formulae of Involatile Substances. — There are, however, 
many chemical substances which are not volatile or which can- 
not be volatilized at temperatures at which we could make 



48 Introduction to General Chemistry 

experimental determinations of their vapor densities. There 
are other solids and liquids which would be decomposed if 
strongly heated. For such substances we could not find chemi- 
cal formulae in the same way as for gases or volatile substances. 
However, we can and do represent by formulae the composition 
of such involatile substances. 

73. The Formula of Red Oxide of Mercury; — The method of 
obtaining the formula of such a substance may be illustrated by 
the case of the red oxide of mercury, which, it will be remem- 
bered, is readily decomposed when heated into mercury and 
oxygen. We find by analysis that this compound contains 
92.6 per cent of mercury and 7.4 per cent of oxygen. By the 
experimental study of volatile mercury compounds, as well as of 
mercury itself, we find that the minimum weight of mercury in 
22.4 liters is 200 g., and therefore represent this weight of 
mercury by the symbol Hg. It now remains to discover what 
multiples of 200 for the mercury and of 16 for the oxygen are 
in the same ratio as the percentages of mercury and oxygen in 
mercury oxide. We find very easily that 200 is to 16 as 92.6 is 
to 7.4, and from this we write the formula HgO. 

We could of course represent the same proportions of mercury 
and oxygen by the formula Hg 2 2 . But we are not able to 
decide which of these to choose as in the case of a volatile sub- 
stance where the formula represents the quantity of material in 
22.4 liters of the gas or vapor under standard conditions. In 
such a case we choose the simpler formula, in this case HgO, but 
we must bear in mind that the formula does not mean quite as 
much in such a case as in that of a gas or volatile substance, 
where it always represents in addition to the true proportion 
of the constituent elements the actual weights of each in 22.4 
liters of the gas under standard conditions. 

74. Symbol Weights and Formula Weights. — The letter or 
pair of letters which represents the minimum weight of an 
element in 22.4 liters of any of its gaseous compounds is called 
the symbol of that element and the weight which this symbol 
represents may then be called the symbol weight. Each of the 
eighty-five or more known elements has been assigned a definite 



Symbols and Chemical Formulae 49 

symbol which represents a definite symbol weight. We have 
seen (62) how the quantities of each element in 22.4 liters of a 
compound gas may b>e represented by a formula made up of 
symbols, each symbol being multiplied by a factor which shows 
how many times the minimum weight of the element is present 
in 22.4 liters of the gaseous compound. The sum of the weights 
represented by the various symbols each multiplied by its factor 
is naturally the weight of 22.4 liters of the gas, represented by 
the formula. This weight is often spoken of as the formula 
weight. In the case of an involatile solid substance the formula 
weight is the weight represented by the formula but indicates 
only theoretically the weight which we should expect 22.4 liters 
of the substance to have if it were a gas under standard con- 
ditions. 

75. The Formulae of Some Elementary Gases. — It is impor- 
tant to note that 22.4 liters of the gases hydrogen, oxygen, 
nitrogen, and chlorine weigh 2, 32, 28, and 71 g. respectively 
(63, Table IV). These weights are for each element just double 
the minimum weights which we find in numerous compounds 
of the elements and therefore in each case just double the weight 
represented by the symbol. We must therefore write, as the 
formulae of these gases, H 2 , 2 , N 2 , and Cl 2 , respectively. The 
formula of an elementary gas in the free state will then represent 
the quantity of that gas in 22.4 liters. We must here point 
out that not every element in the form of gas or vapor is to be 
represented by a formula composed of its symbol taken twice. 
For example, the vapors of mercury and sodium have the single 
symbol formulae Hg and Na, respectively; on the other hand, 
the formulae of the vapors of the elements phosphorus and sulfur 
are P 4 and S 8 . 



CHAPTER VI 

CHEMICAL EQUATIONS 

76. Equations.— In this chapter we shall see how it is possible 
to represent in a very simple way the quantities of substances 
entering into and formed in a chemical reaction. Let us con- 
sider the case of hydrogen and chlorine which has already been 
studied experimentally. We have learned that hydrogen and 
chlorine unite to form hydrogen chloride (44). Furthermore 
we find by experiment that one volume of hydrogen and one 
volume of chlorine give two volumes of hydrogen chloride; so 
that if 22.4 liters of hydrogen united with 22.4 liters of chlorine 
we should obtain 44.8 liters of hydrogen chloride. Now we 
may represent 22.4 liters of hydrogen by the formula H 2 and 
22.4 liters of chlorine by Cl 2 , while for twice 22.4 liters of hydro- 
gen chloride we put the coefficient, 2, in front of the formula and 
write 2HCI. We may then express the facts by stating that H 2 
plus Cl 2 gives 2HCI or 

H 2 +C1 2 ->2HC1 

which may also be written 

H 2 +C1 2 = 2HC1. 

We call either of these expressions the equation for the reaction 
between hydrogen and chlorine. 

77. What an Equation Means. — The equation 

H 2 +C1 2 -» 2 HC1 

expresses the fact that the quantity of hydrogen represented by 
the formula H 2 or 2 g. unites with the quantity of chlorine 
represented by Cl 2 or 71 g. to give the quantity of hydrogen 
chloride represented by 2HCI or 73 g. It also expresses the 
fact that 22.4 liters of hydrogen unite with 22.4 liters of chlorine 
to give 2X22.4 liters of hydrogen chloride, or in general that one 
volume of hydrogen and one volume of chlorine unite to give 
two volumes of hydrogen chloride, the volumes being those of 

50 



Chemical Equations 51 

the gases measured in all cases under standard conditions. In 
reactions involving gases the volume of each gas taken or 
formed is always shown by the coefficient in front of its formula 
in the equation for the reaction. 

78. The Equation for the Burning of Carbon. — Some free 
elements like carbon are not sufficiently volatile to enable us to 
find the formula of the element from measurements of the vapor 
density of the free element, and in such a case we use the symbol 
of the element in writing equations involving its reactions. 
When carbon is burned we find that 1 2 g. of carbon require 3 2 g. 
of oxygen occupying a volume of 22.4 liters, and producing 44 g. 
of carbon dioxide occupying also a volume of 22.4 liters. These 
facts may therefore be represented by the equation 

c+o 2 ->co 2 . 

Here the equation expresses directly the weights of carbon and 
oxygen which unite as well as the weight of carbon dioxide 
formed. At the same time it also shows that 22.4 liters of 
oxygen when completely combined with sufficient carbon gives 
22.4 liters of carbon dioxide, but since the carbon is not in the 
gaseous state the equation does not indicate anything regarding 
the volume of the solid carbon which unites with the volume of 
oxygen represented by the formula 2 . 

79. Solving Problems. — If we remember that the equation 
for the burning of carbon in oxygen is 

C+0 2 ->C0 2 

we may make use of the facts represented by the equation in the 
solution of problems such as the following: How many liters 
of oxygen are required for the burning of 5 g. of carbon? To 
solve this problem we first write down the equation which repre- 
sents the reaction. This shows that the quantity of carbon 
represented by the symbol C, namely, 12 g., requires for its com- 
bustion the volume of oxygen represented by the formula 2 , 
namely, 22.4 liters. Therefore 5 g. of carbon would require 
the volume determined by the proportion 

12:5: :22 .4:x 



52 Introduction to General Chemistry 

where x is the number of liters of oxygen necessary for the com- 
bustion of 5 g. of carbon. In an analogous manner we may 
calculate what volume of carbon dioxide is produced by the 
burning of a known weight of carbon. 

We may also calculate what weight of oxygen is required or 
carbon dioxide produced in the burning of 5 g. of carbon. If 
12 g. of carbon require 32 g. of oxygen, as our equation indicates, 
then we have only to solve the following proportions in order to 
find the weight of oxygen required for 5 g. of carbon : 

12:5:132:3; 

where y is the required answer. 

80. The Burning of Magnesium. — Suppose we desire to find 
by experiment the formula of the product formed by burning 
magnesium in oxygen. It will be recalled that the metal 
magnesium in the form of wire or powder burns with great ease 
in oxygen, forming a white solid substance which we have called 
magnesium oxide (28). We find by experiment that 10 g. of 
magnesium when burned yields 16.6 g. of magnesium oxide. 
Let us suppose that we have discovered by careful experiment 
that magnesium oxide contains only the elements magnesium 
and oxygen. The difference between the weight of the mag- 
nesium oxide formed and the magnesium taken must represent 
the weight of oxygen which has combined with the 10 g. of 
magnesium. This we find to be 6. 6 g. 

Suppose we know that the symbol weight of magnesium is 
24.3 g. or 

Mg = 24.3g. 

It is now required to calculate the relative numbers of symbol 
weights of magnesium and oxygen that unite to form magnesium 
oxide. We know that 10 g. of magnesium unite with 6.6 g. of 
oxygen. We may then make the proportion 

10:6.6: : 24. 3:2 

from which we find that 2=16. Therefore 16 g. of oxygen 
represented by O, combine with the weight of magnesium repre- 



Chemical Equations- 53 

sented by the symbol Mg, and consequently we may represent 
the composition of magnesium oxide by the formula MgO and 
write the equation for the burning of magnesium thus : 

Mg+O^MgO 
or better 

2Mg+0 2 ->2MgO, 

the latter equation having the advantage in that it shows the 
volume of oxygen, 22.4 liters, as well as its weight required for 
the burning of the weight of magnesium represented by 2Mg. 
But since both magnesium and magnesium oxide are solid 
involatile substances the equation does not show the volumes of 
these solids entering into the reaction, as it would in the case of 
gaseous substances. 

81. The Burning of Iron. — It will be recalled (17) that iron 
burns in oxygen, giving iron oxide, the formula for which we 
may now calculate. In an experiment in which 12.6 g. of iron 
was burned the weight of iron oxide produced was 17.4 g., from 
which we find, by subtracting the weight of the iron burned, the 
weight of the oxygen to be 4 . 8 g. These weights of iron and 
oxygen must be in the same ratio that some number of times 
56, the symbol weight of iron, is to some number of times 16 
where these numbers are small integers. Dividing 12.6 by 56 
we get 0.225. Dividing 4.8 by 16 we get 0.300. Since these 
numbers 0.225 and 0.300 are not equal, the formula cannot 
be FeO. It will, however, readily be found that 0.225 i s to 
0.300 as 3 is to 4, and therefore that 12.6:4.8: 13X56:4X16, 
which shows that the formula of the oxide of iron formed by 
burning iron in oxygen is Fe 3 4 . We may then write the equa- 
tion for the burning of iron as follows: 

3 Fe+20 2 ->Fe 3 4 . 

82. The Action of Hydrogen on Copper Oxide. — It will be 
remembered that we found earlier that heated copper oxide and 
hydrogen give metallic copper and water (33) . In a quantitative 
experiment it was found that 2.387 g. of copper oxide yielded 
1.907 g. of copper and 0.54 g. of water. From the weights of 



54 Introduction to General Chemistry 

copper and copper oxide, together with a knowledge of the fact 
that copper oxide is composed of copper and oxygen only, we may 
discover very readily that the formula of copper oxide is CuO, 
knowing the symbol weight of copper to be 63 . 6. Furthermore, 
since water contains only hydrogen and oxygen and o . 54 g. of 
water has been formed from 2.387—1.907 or 0.48 g. of oxygen, 
the weight of hydrogen present in the o . 54 g. of water must have 
been o . 06 g. Making a calculation analogous to that made in 
finding the formula for iron oxide, we find that 0:06:0.48: : 2X 
1:1X16 and that therefore the composition of water is repre- 
sented by the formula H 2 0. We may now write, as the equation 
for the reaction between copper oxide and hydrogen, 

CuO+H 2 ->Cu+H 2 0. 

83. The Action of Acetylene on Copper Oxide. — From what 
has preceded the student will understand that in order to be able 
to write the equation for any reaction we must know all of the 
substances entering into the reaction and all of the products. In 
addition we must know the formula of each substance. We may 
illustrate the method then employed by means of reaction 
between acetylene and copper oxide which we have already 
studied. 

When acetylene is passed over heated copper oxide we obtain 
carbon dioxide and water, while metallic copper is left behind, 
these three substances being the sole products of the reaction (50) . 
The formula of acetylene is C 2 H 2 (62) . The quantity of carbon 
represented by C 2 would give the quantity of carbon dioxide 
represented by 2C0 2 ; and the quantity of hydrogen represented 
by H 2 would give the quantity of water represented by H 2 0, so 
that the quantities of carbon and hydrogen represented by one 
formula weight of acetylene C 2 H 2 would yield the quantities of 
carbon dioxide and water represented by 2C0 2 +H 2 0. The 
quantity of oxygen contained in the quantities of carbon dioxide 
and water represented by 2C0 2 +H 2 is represented by 5O, 
which quantity is contained in the amount of copper oxide 
represented by 5 CuO. It will thus appear that the quantity 
of acetylene represented by C 2 H 2 will require the quantity of 



Chemical Equations 55 

copper oxide represented by 5C11O, and there will be produced 
the quantities of the three products represented by 

2 C0 2 +H 2 0+5Cu. 

The equation is therefore 

C 2 H 2 +5CuO->2C0 2 +H 2 0+5Cu. 

84. The Action of Ammonia on Copper Oxide. — In an analo- 
gous manner we may obtain as the equation for the reaction 
which occurs when ammonia gas is passed over heated copper 
oxide, in which case water, nitrogen, and metallic copper are 

formed, 

2NH 3 + 3 CuO-> 3 H 2 0+3Cu+N 2 . (52) 

85. The Meaning of an Equation. — Since chemists .make 
extensive use of equations, it is of fundamental importance that 
the student should understand exactly how equations are ob- 
tained and what they mean. In every case before the equation 
for the reaction can be written the reaction must have been 
thoroughly investigated by experiment in the manner illustrated 
in the preceding examples. The equation then shows at a glance 
what substances enter into and are formed as a result of the 
reaction. It also shows the composition of each of the substances 
concerned and the proportions in which they take part in the 
reaction, it being assumed in all cases that we know the weight 
for which the symbol of each element stands. 

86. An Equation Balances. — It is one of the most funda- 
mental facts in chemistry that in chemical change no material is 
destroyed but that the elements merely change their forms of 
combination with one another. This important fact, which we 
know as the Law of the Indestructibility of Matter, is also repre- 
sented in every chemical equation. For we notice that in each 
equation we have on each side the same number of symbol 
weights of each element. Thus in the equation 

C 2 H 2 + 5 CuO->2C0 2 +H 2 0+5Cu 

we see that there are on each side two symbol weights of carbon, 
two symbol weights of hydrogen, five symbol weights of copper, 



56 Introduction to General Chemistry 

and five symbol weights of oxygen. This fact is usually expressed 
by saying that the equation balances. 

All of the reactions which we have studied up to this time 
have been thoroughly investigated by chemists and for each the 
reaction equation has been discovered. We may now give, in 
Table VI, a list of such equations for purposes of reference. It 
is not to be expected, however, that the student should make 
great effort to memorize all of these equations, although such a 
task would not be very difficult, for, as a little inspection will 
show, there are certain regularities observable which make this 
a less difficult task than might at first sight seem to be the case. 

TABLE VI 
Equations of Other Reactions Studied 
2Hg+0 2 ->2HgO 
2H 2 +0 2 ->2H 2 
2 Na+ 2H 2 -> 2NaOH+H 2 
NaOH+HCl -> NaCl+ H 2 
2Na+Cl 2 ->2NaCl 
Mg+H 2 0->MgO+H 2 
3Fe+ 4 H 2 0->Fe 3 4 +4H 2 
CH 4 + 2 2 -> C0 2 + 2H 2 
C 3 H 8 + ioCuO -> 3 C0 2 + 4 H 2 0+ 10C11 

87. Problems 

1. What weight of mercury can be obtained by the decom- 
position of 10 g. of mercuric oxide? 

2. What volume of oxygen at o° and 76 cm. can be made 
from 8 g. of mercuric oxide? 

3. What weight of sodium must be acted on by water to 
yield 500 c.c. of hydrogen at o° and 76 cm.? 

4. What weight of common salt can be made from 10 g. of 
metallic sodium? 

5. What volume of hydrogen at 20 and 72 cm. would be 
formed by the action of sufficient steam on 6 g. of magnesium? 

6. What weight of copper oxide would be required for the 
oxidation of 200 c.c. of propane measured at 25 and 74 cm.? 
(See last equation of Table VI above.) 

What weight of water would be formed? 



CHAPTER VII 
ACIDS, BASES, AND SALTS— I 

88. Caustic Soda or Sodium Hydroxide. — Let us now con- 
sider the chemical changes which occurred in the formation of 
common salt from metallic sodium, which we have already 
studied experimentally. It will be recalled that sodium reacted 
violently with water, giving hydrogen and sodium hydroxide, the 
reaction being represented by the equation 

2 Na+ 2 H 2 -> 2 NaOH+H 2 . (40) 

If we repeat the experiment and evaporate the water we find 
that sodium hydroxide (also known as caustic soda) is left as a 
white solid which is readily soluble in water. This solution feels 
" soapy" to the fingers and if greatly diluted with water is found 
to have an unpleasant "soapy" taste. (It must not be tasted 
unless greatly diluted with water, since the concentrated solu- 
tion acts powerfully on the mucous membrane.) A piece of red 
litmus paper is turned blue if dipped in the solution. We know 
many other substances which have properties similar to those 
of sodium hydroxide. Such substances are called bases; they 
also have other characteristic properties, the most important of 
which we may now consider. 

89. Bases Neutralize Acids. — We have learned (41) that 
caustic soda and hydrochloric acid (which is a solution of hydro- 
gen chloride in water) react to give common salt. The equation 
for this reaction is 

NaOH+HCl-> H 2 0+NaCl. 

If we add more than sufficient of the acid and then evaporate 
the solution to dryness, the excess of hydrogen chloride will pass 
off with the water and nothing but pure salt, the chemical name 
of which is sodium chloride, will remain. If we test hydro- 
chloric acid with blue litmus we find that the latter is turned red, 
even by a very dilute solution. But we find that a solution of 

57 



58 Introduction to General Chemistry 

pure common salt in water has no effect on either blue or red litmus: 
it is neutral. 

90. Properties of Acids. — If we again add, drop by drop, a 
solution of hydrogen chloride to one of sodium hydroxide to 
which a few drops of a solution of litmus have been added, we 
find that the change of color from blue to red is produced suddenly 
and not gradually, a single drop being sufficient to cause the 
change. If we stop adding hydrogen chloride at this point we 
find that the solution consists only of pure salt and water (with 
but a minute amount of litmus). It no longer has the taste of 
the sodium hydroxide, but only that of salty water. A diluted 
solution of hydrogen chloride has a rather agreeable sour taste, 
reminding one of vinegar or lemon juice. Our experiment has 
shown that both the taste and the behavior toward litmus of 
sodium hydroxide and hydrogen chloride have been changed in 
their interaction. We say that they have neutralized each other. 
We know very many substances which will neutralize sodium 
hydroxide; all of these have a sour taste and color litmus red. 
We call such substances acids, the common name of hydrogen 
chloride solution being hydrochloric acid. 

91. Another Base; Ammonium Hydroxide. — As we have 
already seen (51), ammonia gas dissolves readily in water, giving 
a solution which turns litmus blue, and we are not surprised to 
find that it neutralizes hydrochloric acid. If we evaporate the 
neutralized solution we obtain a white crystalline substance, the 
composition of which is represented by the formula NH 4 C1. 
Since ammonia gas has the formula NH 3 and hydrogen chloride 
the formula HC1, we might be inclined to write the equation 

NH 3 +HC1->NH 4 C1, 

and, in fact, just this reaction takes place if we bring the two 
gases together, a dense white cloud of the solid product being 
formed. However, if a very concentrated solution of ammonia 
in water is cooled to a very low temperature, we may obtain 
crystals of a substance having a composition represented by the 
formula NH 4 OH and called ammonium hydroxide. This sub- 
stance is formed thus: 

NH 3 +H 2 0->NH 4 OH. 



Acids, Bases, and Salts — / 59 

We might think to obtain it by the evaporation of the water 
solution of ammonia ; but instead we get only ammonia gas and 
water vapor. In fact, the crystals of ammonium hydroxide 
obtained at a low temperature undergo a similar change if they 
are not kept very cold. We say that ammonium hydroxide 
dissociates readily into ammonia and water. Chemists think that 
in a water solution of ammonia part of the latter is combined 
with water to form ammonium hydroxide. It is this substance 
which is thought to act directly on red litmus, changing it to 
blue, and to act on hydrochloric acid as follows: 

NH 4 OH+HCl -> NH 4 C1+H 2 0. 

We therefore call ammonium hydroxide a base. 

92. Ammonium Chloride, Salts. — The substance NH 4 C1 is 
called ammonium chloride. In appearance, taste, and other 
properties to be studied later, sodium chloride and ammonium 
chloride closely resemble one another. They are examples of an 
important class of chemical substances called salts. 

A review of the two neutralizations just discussed will show 
that they have much in common : in each case a base reacts with 
an acid to form a salt and water. Somewhat later, other important 
facts regarding neutralization will be discovered. Before dis- 
cussing such matters we will first become acquainted with a few 
other important acids, bases, and salts. 

93. Sulfuric Acid. — One of the most important, if not the 
most important, of all acids is a substance which is known as oil 
of vitriol or sulfuric acid. It is manufactured in immense 
quantities and is very cheap, the commercial grade selling for 
less than one cent a pound. We shall not now consider the 
method of its manufacture further than to state that it is made 
from sulfur. Its composition is represented by the formula 
H 2 S0 4 . It is a colorless liquid of "oily" consistency, but is 
not really an oil, as it will mix with water in all proportions. 
It must be handled with caution, since it can cause bad burns 
if it is spilled on the skin. {In case of accident, wash of the acid in 
much running water, immediately.) When sulfuric acid is mixed 
with water, the mixture gets boiling hot, for which reason the acid 



60 Introduction to General Chemistry 

should be added very slowly, with stirring, to the water, if a 
dilute solution is to be made. 

94. Neutralization of Sulfuric Acid, Sodium Sulfate. — We 
find that the dilute solution has a sour taste and that it turns 
litmus red. We may next try whether it will neutralize a solu- 
tion of sodium hydroxide, for which purpose we may add to a 
dilute solution of sulfuric acid a few drops of litmus solution and 
then run in sodium hydroxide solution drop by drop until neutral- 
ity is reached. If the neutral solution is now boiled until a solid 
begins to appear and then is left to evaporate at room tempera- 
ture, large, transparent, glassy-looking crystals will be formed. 
These crystals dissolve readily in water to form a neutral solu- 
tion, which does not have a sour taste. 

If we allow the dry crystals to remain in the open air we find 
that they lose weight rapidly and turn white upon the surface, 
forming a fine white powder. Finally nothing is left of the large, 
clear, glassy crystals; only the powder remains, the weight of 
which is much less than that of the original material. What is 
the cause of this curious change? Let us put one of the large 
clear crystals into a dry test tube and heat gently the lower end 
of the tube containing the crystal, while the tube is held nearly 
horizontally. We soon see that water has collected in large 
amount in the cold end of the tube, while only a white powder 
is left behind. It is now easy to understand what occurred when 
the large crystal was exposed in the open air. It dissociated into 
the white powder and water which disappeared as vapor. The 
analysis of the thoroughly dried powder would show that it 
contains only sodium, sulfur, and oxygen, and in the proportions 
represented by Na 2 S0 4 , and since the clear crystals yielded only 
Na 2 S0 4 and water, their composition must be represented by 
Na 2 S0 4 .^H 2 0, where x is a whole number which must be found 
by means of a quantitative analysis. We call the original sub- 
stance the hydrate of sodium sulfate, a hydrate of a salt being 
a compound of the salt with water. 

We may now make the equation for the formation of this 
salt from sulfuric acid. We took H 2 S0 4 and NaOH and got 
Na 2 S0 4 , from which we see that if two formula weights of water 



Acids, Bases, and Salts — / 61 

were formed from one formula weight of H 2 S0 4 and two of NaOH, 
the whole of the material taken would be accounted for thus : 
H 2 S0 4 + 2NaOH -> Na 2 S0 4 + 2H 2 0. 

This conclusion is rendered probable by the fact that in the other 
neutralizations we have studied water was always one of the 
products; it may be confirmed by mixing with dry sodium 
hydroxide pure sulfuric acid, whereupon water and Na 2 S0 4 will 
result. The salt Na 2 S0 4 is called sodium sulfate. Crystals of 
anhydrous sodium sulfate are different in form from those of the 
hydrate. 

95. Quantitative Analysis of a Hydrate. — Let us now consider 
the quantitative composition of the large, glassy crystals which 
yielded Na 2 S0 4 and water. If we weigh a crystal contained in a 
porcelain dish and allow it to stand a day or two at room tempera- 
ture we find that only the white powder remains. If we now heat 
the dish and contents over a flame in order thoroughly to dry 
the powder, and let it cool and weigh it again, it is obvious that 
the loss of weight will represent the weight of water originally 
combined with the weight of dry Na 2 S0 4 left in the dish. 

96. Sodium Sulfate Decahydrate: Na 2 S0 4 «ioH 2 0. — Now 
suppose that 5 . 796 g. of the hydrate of sodium sulfate yielded 
2.556 g. of dried sodium sulfate, Na 2 S0 4 , what is the formula of 
the hydrate? In other words, what is the numerical value of 
x in the formula Na 2 S0 4 *xH 2 0? The weight of water driven off 
was 5.796— 2.556 = 3.240 g. We may therefore write the 
proportion, 2.556 is to 3.240 as the formula weight of sodium 
sulfate is to the x times the formula weight of water. Now, the 
formula weight of sodium sulfate is 2X23+32+4X16 = 142 
and that of water is 2X1 + 16 = 18. Therefore 2.556:3.240:: 
142 : i8#, from which we find that x = 10, and are thus led to the 
conclusion that the hydrate of sodium sulfate has the formula 
Na 2 S0 4 «ioH 2 0. If the reaction between sulfuric acid and 
sodium hydroxide is represented by the equation H 2 S0 4 + 
2NaOH->Na 2 S0 4 +2H 2 0, then the hydrate Na 2 S0 4 -ioH 2 
must have resulted from the union of the sodium sulfate with 
part of the water which formed the solution, thus : 

Na 2 S0 4 + ioH 2 -> Na 2 S0 4 . ioH 2 0. 



62 Introduction to General Chemistry 

This substance is called sodium sulfate decahydrate [deca mean- 
ing ten). 

97. Hydrates. — Sodium sulfate forms other compounds 
with water, namely Na 2 S0 4 «7H 2 and Na 2 S0 4 -H 2 0; but the 
decahydrate is the common one. Many other salts form 
hydrates and some form a series of hydrates, as this salt does. 
But it must not be supposed that all salts form hydrates. For 
example, sodium chloride and ammonium chloride do not. 

Solutions of the hydrated salt have exactly the same prop- 
erties as those of solutions of the anhydrous salt. 

98. Sodium Hydrogen Sulfate: NaHS0 4 . — If we exactly 
neutralize a definite quantity of sulfuric acid with a solution of 
sodium hydroxide, noting the volume of the latter used, and again 
add to a second portion of sulfuric acid, exactly equal to the first, 
exactly half as much sodium hydroxide solution as that used 
in the first case, we find that the first solution yields when 
evaporated pure sodium sulfate, Na 2 S0 4 ; while the second gives 
crystals having a different shape and appearance, and different 
chemical properties. Analysis shows that the composition of 
these crystals is represented by the formula NaHS0 4 . The 
substance is called sodium hydrogen sulfate. The equation for 
the reaction in the second case is 

H 2 S0 4 +NaOH-> NaHS0 4 +H 2 0. 

99. The Law of Definite Composition Again. — We may now 

consider one of the most important and fundamental of all 
chemical questions, namely, whether the proportions of the 
elementary constituents of a substance are dependent upon the 
proportions which we take to the substances from which we form 
the substance in question. For example, we may inquire 
whether we could get a sulfate of sodium with a somewhat 
larger or smaller percentage of sodium if we had used, in the 
preceding experiment, other proportions of acid and base. 
Experiment will show, however, that if we had added a little 
more or less sodium hydroxide we would still have been able to 
obtain much NaHS0 4 , but that in such cases there would also 
be some Na 2 S0 4 formed or a little free sulfuric acid left after all 



Acids, Bases, and Salts — / 6$ 

the NaHS0 4 had been separated from the water. Facts like 
these which are met with on every hand give a special significance 
to the Law of Definite Composition. 

ioo. Acid Properties of Sodium Hydrogen Sulfate. — We see 
that sodium sulfate, Na 2 S0 4 , contains exactly twice the weight of 
sodium for a given weight of sulfur and oxygen as does sodium 
hydrogen sulfate, NaHS0 4 . Moreover, we have become 
acquainted with the important fact that sulfuric acid can form 
two sorts of sodium salts. If we dissolve crystals of sodium hydro- 
gen sulfate in water, we find that the dilute solution has a sour 
taste and it turns litmus red, for which reasons we should be 
inclined to say that it has acid properties. In accord with this 
view, we find that the solution will readily neutralize a solution 
of sodium hydroxide, giving sodium sulfate and water, thus: 

NaHS0 4 +NaOH -> Na 2 S0 4 +H 2 0. 

101. Ammonium Sulfate and Ammonium Hydrogen Sul- 
fate. — If we completely neutralize sulfuric acid with a solution 
of ammonium hydroxide, we obtain a salt called ammonium sul- 
fate (NH 4 ) 2 S0 4 , thus: 

H 2 S0 4 + 2 NH 4 OH -> (NH 4 ) 2 S0 4 -f- 2H 2 ; 

while with half the proportion of ammonium hydroxide we obtain 
ammonium hydrogen sulfate, thus : | 

H 2 S0 4 +NH 4 OH -> NH 4 HS0 4 +H 2 0. 

102. Monobasic and Dibasic Acids: Acid Salts and Neutral 
Salts. — Hydrochloric acid reacts with sodium hydroxide only in 
one proportion, thus: 

HCl+NaOH -> NaCl+H 2 0, 

for which reason we call it a monobasic acid; but since one 
formula weight of sulfuric can Unite with a maximum of two 
formula weights of sodium hydroxide we call sulfuric acid a 
dibasic acid. Salts in which but half the maximum quantity 
of base has been neutralized are usually called acid salts, because 
they still have acid properties. Thus we frequently speak of 
sodium acid sulphate, meaning NaHS0 4 . Chemists know many 



64 



Introduction to General Chemistry 



other dibasic acids, all of which also can form acid salts as well as 
neutral salts, as salts like Na 2 S0 4 are called. 

103. Making Hydrochloric Acid from Common Salt. — If we 
place in a flask (Fig. 23) 58 g. of dry common salt and 100 g. of 
sulfuric acid, to which 30 g. of water have been added, and warm 
the mixture, a change occurs with the production of a colorless 
gas which dissolves in water very readily, giving a solution which 
we can easily recognize as hydrochloric acid. After the action 
of the sulfuric acid on the salt is complete, a white solid is left in 
the flask, which may easily be dissolved in water. By evaporat- 
ingpart of the water, and letting the solution stand a while, we 

may obtain colorless, transparent 
crystals of sodium hydrogen 
sulfate. The following equation 
represents the reaction: 

NaCl+H 2 S0 4 -> NaHS0 4 +HCl. 

We have to deal here with a new 
__ sort of chemical change — one in 
which an acid acts upon a salt of 
another acid to give a salt of the 

first acid and to produce the acid corresponding to the first salt. 

This is a very important kind of chemical reaction, which we 

shall frequently make use of, since by its means we may make 

acids from their salts. 

104. Making Nitric Acid from Chile Saltpeter. — We shall now 
use the method just described for the preparation of a new acid 
from a white, crystalline substance called Chile saltpeter, which 
is found in large quantities as a mineral substance in the desert 
region of Chile. 

If we place 85 g. of Chile saltpeter in a retort (Fig. 24), add 
100 g. of sulfuric acid, mixed with 30 c.c. of water, and then heat 
the mixture gently, a yellow-colored liquid may be collected in a 
cooled flask. This yellow liquid gives off a brown gas and 
becomes colorless when boiled a few minutes. Its analysis shows 
its formula to be HN0 3 and it is called nitric acid. It is a color- 
less liquid which may be boiled and distilled in glass vessels. 




Fig. 23 



Acids, Bases, and Salts — / 



65 



Pure or concentrated nitric acid is even more dangerous than 
sulfuric acid, causing serious burns and destroying clothing, 
and must be handled with greatest care. It will mix with water 
in all proportions, giving a solution which, when very dilute, has 
a sour taste and turns litmus red. 

When nitric acid is mixed with sodium hydroxide solution the 
latter is neutralized, a salt of the composition NaN0 3 and water 
being the only products, as represented by the equation 

HN0 3 +NaOH -> NaN0 3 +H 2 0. 

The salt, which is called sodium nitrate, is found to be identical 

with purified Chile saltpeter. The 

action of sulfuric acid on saltpeter 

leaves in the retort a white solid 

which closely resembles that left 

when salt is heated with sulfuric acid, 

and, in fact, the residue is easily found 

to be the same substance, sodium 

hydrogen sulfate, NaHS0 4 . The 

equation for the reaction is therefore 

NaN0 3 +H 2 S0 4 -> NaHS0 4 +HN0 3 




Fig. 24 



105. The Action of Nitric Acid on Ammonium Hydroxide. — 

We may now propose a question to be answered, not after direct 

experiment, but as a result of the general knowledge we have 

gained regarding the behavior of the acids and bases already 

studied. It is: What would be the result of mixing nitric acid 

and ammonium hydroxide? We recall that hydrochloric acid 

and sodium hydroxide, a base, give sodium chloride and water, 

thus: 

HCl+NaOH -> NaCl+H 2 ; 

that the same acid gives with ammonium hydroxide, also a base, 
ammonium chloride and water, thus : 

HCl+NH 4 OH -> NH 4 C1+H 2 0. 

Furthermore, we have just seen (104) that nitric acid and sodium 
hydroxide give sodium nitrate and water, thus : 
HN0 3 +NaOH -> NaN0 3 +H 2 0, 



66 Introduction to General Chemistry 

and we would certainly expect that nitric acid and ammonium 
hydroxide would behave analogously and give ammonium nitrate 
and water, thus : 

HN0 3 +NH 4 OH -> NH 4 N0 3 +H 2 0. 

Now this is precisely what takes place when we test our prediction 
by experiment. We seem, therefore, to have discovered the 
secret of the way in which acids and bases act toward each other. 
It may be summed up in the statement, An acid and a base 
neutralize each other, forming a salt and water. 

106. A New Base: Caustic Potash or Potassium Hydroxide. 
— Let us now take up the study of a new base, caustic potash, 
which closely resembles caustic soda (sodium hydroxide). It 
will be remembered that the metal sodium reacts violently with 
water, giving sodium hydroxide and hydrogen gas, thus: 

2 Na+ 2 H 2 -> 2 NaOH+H 2 (88) 

Now, chemists know another metallic element, potassium, which 
closely resembles sodium. Like sodium, it is a silver-white 
metal, soft enough to be cut easily with a knife and tarnishing 
very rapidly in the air. For a reason that we shall soon learn it 
is kept covered with oil in a carefully stoppered bottle. If we 
throw a small bit of potassium into a beaker of water, it bursts 
into a flame of lavender color, spinning and darting to and fro 
on the surface of the water and completely disappearing in a few 
moments. Examination of the water shows that it will turn 
litmus blue, that it has a "soapy" taste, like a very dilute solu- 
tion of sodium hydroxide, and that a white solid is left when the 
solution is evaporated to dryness. This solid is found by suitable 
methods of analysis to contain the elements potassium, oxygen, 
and hydrogen in the proportion represented by the formula 
KOH, and is called potassium hydroxide. 

If instead of throwing the bit of potassium on the surface of 
the water we bring it under the mouth of an inverted cylinder 
filled with water, with the mouth immersed in a vessel of water, 
the potassium rises to the top of the water in the cylinder, pro- 
ducing a gas which displaces the water in the cylinder, but does 



Acids, Bases, and Salts — / 67 

not take fire. The gas is easily identified as hydrogen, while the 
water contains dissolved potassium hydroxide as before. The 
equation for the reaction in the cylinder is 

2 K+2H 2 0->2KOH+H 2 . 

When the action takes place in the open beaker, the heat pro- 
duced sets fire to the hydrogen, which burns, together with a small 
portion of the potassium. 

107. Potassium Salts. — On account of the behavior of a solu- 
tion of potassium hydroxide toward litmus and also because of its 
" soapy" feel and taste, we should conclude that it is a base and 
if so that it should form salts with acids. We might even venture 
to predict the formulae of the salts it would be expected to form 
with hydrochloric, sulfuric, and nitric acids, and to write the 
equations as follows: 

HCl+KOH-> KC1+H 2 
H 2 S0 4 + 2KOH -> K 2 S0 4 + 2 H 2 
H 2 S0 4 +KOH -> KHS0 4 +H 2 
HNO3+KOH -> KN0 3 +H 2 0. 

And in every case these predictions would be found by experiment 
to be correct! The potassium salts so formed are all white 
crystalline solids and are all soluble in water. All except potas- 
sium hydrogen sulfate give solutions which are neutral to litmus, 
while this salt has acid properties like those of sodium hydrogen 
sulfate. 



CHAPTER VIII 
WATER AND SOLUTIONS 

1 08. Water. — We have already learned that pure water is 
readily obtained by the distillation of natural waters (23), and 
that it is a compound of hydrogen and oxygen, the composition 
of which is represented by the formula H 2 (70) . In describing 
a substance we shall often mention its physical and chemical 
properties. The properties of a substance embrace: the state 
(whether solid, liquid, or gaseous); crystalline form, if solid; 
specific gravity or density; color; odor; taste; conductivity for 
heat and electricity; boiling-point; freezing-point, etc. The 
chemical properties of a substance are those which it exhibits in 
its typical chemical reactions. 

109. The Physical Properties of Water: Color. — We know 
that according to the temperature water can exist as solid, liquid, 
or gas. The color of liquid water is a very faint blue; so faint, in 
fact, that it cannot be noticed in a glass of water, but is obvious 
in a white, bathtub full of clear water. The color of large bodies 
of clear water is usually blue, but it may be of any other shade 
if dissolved or suspended impurities (mud) are present. The 
yellow color of the waters of many rivers is due to suspended 
clay; such water is not clear, but muddy or turbid. Streams and 
lakes in hemlock forests often contain perfectly clear water 
having the color of tea, due to coloring-matter dissolved from 
the hemlock. The clear green color of some waters is usually 
the result of the blending of the natural blue color of the water 
with the yellow light reflected from the sand beneath. 

no. Specific Gravity or Density. — At the temperature of 
4 C, 1 c.c. of water weighs 1 g. Since the specific gravity or 
density of any substance may be defined as the weight of 1 c.c, it 
follows that water has a specific gravity of 1 . 000 at 4 C. Or, we 
may say that the specific gravity or density of a substance is 
found by dividing its weight by the weight of an equal volume of 
water. Water has its greatest density at 4 ; if a given volume 

68 



Water and Solutions 69 

of water at 4 is either heated or cooled, it expands and therefore 
decreases in density. 

in. Specific Heat. — The quantity of heat required to raise 
the temperature of 1 g. of water i°C. is by definition called one 
calorie. Water is said to have a specific heat of one or unity. 
The specific heat of any substance is the quantity of heat in calories 
required to raise the temperature of one gram of it one degree. 
Nearly all substances have specific heats less than unity. 

112. Vapor Pressure. — Water contained in an open vessel 
evaporates at all temperatures, but the more rapidly in propor- 
tion as the temperature is higher, other things being equal. If 
water evaporates into an evacuated space the 
pressure within the space increases to a value 
which is dependent only upon the tempera- 
ture, being greater in proportion as the 
temperature is higher. The pressure so pro- 
duced is called the vapor pressure of water; 
it may easily be demonstrated by means of a 
barometer tube filled with mercury. If we 
prepare two such tubes (Fig. 25) and intro- 
duce a few drops of water into one by means 
of a suitably shaped glass tube, the water fig. 2S 

will rise until it floats on the surface of the 
mercury. At the same time the level of the mercury will 
fall 2 or 3 cm., showing that a pressure has been produced 
above the mercury in the space which has been a vacuum. 
If the tube into which the water is introduced has a glass jacket 
into which warm water can be poured, it will be found that 
the higher the temperature is, the higher the vapor pressure will be. 
If we should raise the temperature to ioo°, the level of the 
mercury in the barometer tube would sink to that of the surface 
of the mercury in the dish in which the tube stands, thus show- 
ing that the vapor pressure at ioo° is equal to the pressure of the 
atmosphere. Table VII shows the vapor pressure of water at 
various temperatures between o° and ioo°. 

When the atmospheric pressure is 760 mm., water boils at 
ioo°. Now, we see from the table that at ioo° the vapor 



JL 



7° 



Introduction to General Chemistry 



pressure is 760 mm., therefore the boiling-point is that temperature 
at which the vapor pressure becomes just equal to the normal atmos- 
pheric pressure, 760 mm., which is the average pressure at sea- 



TABLE VII 



Temperature 


Pressure 


Temperature 


Pressure 


o° 

10 

20 

30 

40 

50 


4.6 mm. 
9.2 

17-4 
31.6 
55-o 
92.2 


6o° 

7o 

80 

90 

99 

100 


149 . 2 mm. 

233-8 

355-5 

526.0 

733-2 

760.0 



level. At higher altitudes, at which the atmospheric pressure 
is less than 760 mm., water boils at temperatures lower than 
ibo°. Thus if the pressure is 733.2mm., the boiling-point is 
99 / Since the atmospheric pressure at a given place is variable 
through a range of 20 mm. or more, the boiling-point at this 
place is not constant, but varies with the rise and fall of the 
barometer. 

1 13. Correction of the Volume of a Gas for Vapor Pressure. — 
Gases like hydrogen and oxygen, which are not very soluble in 
water, are often measured in tubes in which the 
gases are confined by means of water. Such gases 
always contain water vapor, and part of the total 
pressure exerted by the gas is due to the vapor 
pressure of the water. The part of the pressure 
(partial pressure) exerted by the gas itself is found 
by subtracting from the total pressure , the vapor 
pressure of the water. For example, suppose that 
some hydrogen is collected over water in a grad- 
uated glass tube (Fig. 26). If the position of the 
tube is adjusted so that the level of the water is the same inside 
the tube as outside, the total pressure within must be exactly 
equal to the atmospheric pressure, as shown by the barometer. 
Suppose that the barometric pressure is 748 . 6 mm. and the 
temperature 20 . Table VII shows that at 20 the vapor 
pressure of water is 17.4 mm., therefore the pressure due to the 



Fig. 26 



Water and Solutions 71 

hydrogen is 748 .6— 17. 4 = 73 1.2 mm. If the observed volume 
was 30 ex., the volume, V, at standard conditions would be 

30X731^2X273 . 

V = t-—z = 20.0 C.C. 

760X293 

114. Vapor Pressure of Liquids and Solids in General. — 

Liquids in general readily pass into the form of vapor, and just 
as in the case of water, a given pure liquid has, at each tempera- 
ture, a definite vapor pressure; but the vapor pressure of one 
liquid — say alcohol — is not in general the same at a given 
temperature as that of another liquid — say water. In every 
case, however, the boiling-point of the liquid is that temperature at 
which its vapor pressure equals 760 mm. Many solids, for 
example, camphor and naphthalene (moth-balls), have appre- 
ciable vapor pressures at room temperature; but the vapor 
pressures of most solids at such temperature are too small to be 
noticeable. 

115. Latent Heat of Evaporation. — If it is true that water 
boils at ioo° because at this temperature the vapor pressure of 
water just equals the normal atmospheric pressure, it may be 
asked why the whole of the water does not change at once into 
steam as soon as its temperature is raised to ioo°. We know, of 
course, that this does not occur, and, further, that the rapidity 
with which water boils away is greater, the greater the amount 
of heat applied. The explanation is found in the fact that it 
requires a large amount of heat to change water at ioo° into steam 
at the same temperature. In fact, 540 calories of heat are required 
for the conversion of 1 g. of water at ioo° into steam. The heat 
so used up does not raise the temperature of the substance. It is 
consumed in changing the liquid water into the gaseous state; 
it is said to become latent, and in consequence we say that the 
latent heat of evaporation of water is 100 calories. Every pure 
liquid has a latent heat of evaporation. This differs from one 
substance to another. 

116. Use of Steam for Heating. — When steam cools to ioo° it 
begins to condense to liquid water, and for every gram of steam 
that condenses 540 calories of heat are given out. The heat 



72 Introduction to General Chemistry 

so given out may be considered to be that which became latent 
when the water was, by being heated, converted into steam. It 
is on account of the latent heat given out upon condensation that 
steam is so effective in the heating of buildings: every gram of 
steam that condenses in the radiator liberates 340 calories of 
heat. Of course, the further cooling of the water in the radiator 
gives out some additional heat. 

117. Burns Produced by Steam. — It is a well-known fact 
that serious burns result when steam comes in contact with the 
skin. At first thought, this result seems to be out of harmony 
with the fact that air at ioo° can be borne by the hand without 
discomfort. The explanation of this difference is found in the 
fact that gases (including the vapors of boiling liquids) are very 
poor conductors of heat as compared with liquids. Steam at ioo° 
partly condenses on striking the skin and wets it with a layer 
of boiling-hot water, which is a good conductor of heat. Further- 
more, since 540 calories of heat are given out by every gram of 
steam condensed to water, the latter is kept at ioo° as long as 
steam is present. On the other hand, air is so poor a conductor 
of heat that the skin is not burned by a brief exposure to it at ioo°. 

118. Latent Heat of Fusion of Ice. — Ice melts at o°; but all 
of a given mass of ice does not melt immediately when its 
temperature is raised to zero. Just as heat is required to change 
liquid water into vapor, so also heat is needed to change ice at 
zero into water at the same temperature. The heat so absorbed 
is called the latent heat of fusion of ice. It requires 79 calories 
to melt 1 g. of ice; therefore the latent heat of fusion of ice is 
79 calories. Every solid has a definite and characteristic latent 
heat of fusion. 

119. The Density of Ice. — The density or specific gravity of 
ice is 0.917. It is for this reason that ice floats on water. The 
expansion which occurs when water freezes exerts very great 
pressure, illustrations of which are often seen in the bursting of 
water pipes and other vessels when water freezes in them. Not 
all liquids expand upon freezing; in many cases contraction 
occurs, thereby giving rise to solids which sink in the correspond- 
ing liquids. 



Water and Solutions 73 

, 120. Solutions and Suspensions. — The mixture which results 
upon dissolving salt in water is called a solution of salt in water. 
The terms " dissolve" and "solution" are used in chemistry with 
definite meanings. If, upon mixing a solid with a liquid, the 
former partly or wholly disappears and the resulting liquid is 
still clear and transparent and not cloudy or muddy, and if, 
moreover, upon allowing the liquid to evaporate we regain the 
unchanged solid substance, we say that the solid had dissolved 
in the liquid to form a solution. Either or both of the substances 
may be colored and still a clear (although colored) solution may 
result. The liquid in which a substance is dissolved is called the 
solvent. 

If we stir up some common clay with water, much of the clay 
fails to settle out of the water at once, and we get a cloudy or 
muddy fluid, like the water of a muddy river. In this case we 
do not say that the clay has dissolved in the water or that we have 
a true solution of the clay. We say that the clay is suspended 
in the water, and call the muddy water a suspension. Clay 
suspended in water will settle out very slowly and finally leave 
clear water above a layer of mud. 

121. The Concentration of Solutions. — A solution containing 
a small proportion of a dissolved substance is said to be dilute, 
while one containing a large proportion is called concentrated. 
We dilute a concentrated solution by adding solvent to it, and 
concentrate a dilute solution by evaporating the solvent. We use 
the term concentration in discussing the relative amount of 
dissolved substances in a solution. 

122. Solubility of Substances: Saturated Solutions. — It is 
easy to discover that the amount of a substance which will dis- 
solve in a given amount of water, say 100 c.c, depends upon the 
nature of the substance and upon the temperature. If we mix 
some common salt with about double its weight of water and 
stir or shake the mixture a sufficient length of time (usually one 
to two hours), keeping the temperature constant all the while, 
and then, after allowing any suspended crystals to settle, draw 
off a portion of the clear solution, weigh it, and evaporate the 
water, we get the salt dissolved in the portion of the solution 



74 Introduction to General Chemistry 

taken. By weighing the salt we can readily find the weight Qf 
salt dissolved in a given weight of water at the temperature at 
which the experiment was made. We find in this way that 
ioo g. of water at 25 dissolves 37.6 g. of salt. 

To make such a solubility determination we must observe 
several precautions: First, the amount of solid substance must 
be considerably greater than the amount of water taken will dis- 
solve; secondly, the shaking must be continued as long as more 
substance dissolves — this is easily ascertained by prolonging the 
shaking and making additional determinations of the concentra- 
tion of the solution; thirdly, the temperature must be kept 
constant. 

A solution which at a fixed temperature will dissolve no more 
of a given substance is called a saturated solution. When we 
speak of the solubility of a substance we mean the amount of 
substance dissolved in a given amount of water in the case of 
the saturated solution. The following brief table gives the 
solubilities in water at 25 of several salts. 

TABLE VIII 

Grams of Substance in ioo g. of Water at 25 



123. Supersaturated Solutions. — At 25 100 g. of water will 
dissolve 27 g. of sodium sulfate decahydrate, Na 2 S0 4 'ioH 2 0, 
while at 30 the same amount of water will dissolve 40 g. of the 
salt. If we make a saturated solution of the salt at 30 , having 
an excess of crystals of the salt present, and then cool the whole 
to 2 5 , and keep it at 25 , stirring or shaking it for an hour or two, 
more solid is deposited and there results a solution which contains 
just the same weight of the salt in 100 g. of water as a saturated solu- 
tion at 2 5 , namely, 27 g. 

A slight change in the procedure gives a very different result 
and brings to light a new phenomenon. If the solution of sodium 
sulfate which is saturated at 30 is freed from every particle of the 



Water and Solutions 75 

solid crystalline substance and then allowed to cool to 25 or even 
lower, without being stirred or shaken, it remains perfectly clear and 
does not deposit any crystals. Such a solution contains at 25 
much more sodium sulfate than a saturated solution prepared 
at 2 5 in the manner described in the preceding paragraph. 
This more concentrated solution is called a supersaturated 
solution. If we now drop into the supersaturated solution a 
crystal of sodium sulfate (and for this purpose an almost in- 
visible fragment of the crystalline dust will be sufficient), the 
formation of crystals will begin at once and proceed until the 
amount of dissolved substance per 100 g. of water is reduced 
exactly to that of a saturated solution at the existing tempera- 
ture. 

Experience has shown that a supersaturated solution can 
only be obtained in the complete absence of the solid substance, 
and that a supersaturated solution begins to deposit its excess 
of dissolved substance when a crystal of this same substance is 
brought into the solution. The deposition of crystals by a 
supersaturated solution can also often be started by shaking 
or stirring the solution or by adding a crystal of another sub- 
stance having the same crystalline form. 

Not all substances form supersaturated solutions equally 
readily. The presence of impurities favors supersaturation. 
Syrups, preserves, and honey are often supersaturated with 
respect to the sugar dissolved in the water present. When such 
solutions "turn to sugar," this is only the crystallization of the 
excess of sugar above that required to make a saturated solution. 

124. Solubility of Liquids in Liquids. — It is proverbial that 
"oil and water will not mix." On the other hand, some pairs of 
liquids will mix completely in all proportions; examples of such 
combinations are water and alcohol and water and sulfuric acid. 
We know other pairs of liquids that will not dissolve one another 
in all proportions, but that will dissolve one another partially. 
Water and ether belong to this class; 100 c.c. of water will 
dissolve 8 c.c. of ether, and 100 c.c. of ether will dissolve 3 c.c. 
of water. If we pour ether into water, we find that the former 
floats on the surface of the latter. If equal volumes of ether and 



76 Introduction to General Chemistry 

water are thoroughly shaken together, the former soon separates 
from the latter, and two distinct layers result as before. If, now, 
we examine each layer, we find that the water contains some 
dissolved ether and the ether some dissolved water. This is a 
case of partial miscibility. 

125. Solubility of Gases in Liquids. — We have already 
learned that hydrogen chloride (44) and ammonia (51) are both 
very soluble in water. At o° water dissolves 550 times its own 
volume of the first gas and 1,150 times its volume of the second. 
No gas which we have studied is completely insoluble in water; 
for example, 100 c.c. of water dissolves 2 . 1 c.c. of hydrogen and 
4.8 c.c. of oxygen. Fishes depend for their existence upon the 
oxygen dissolved in water; by means of their gills they take 
from the water the oxygen they require. 

126. Henry's Law. — The solubility of all gases decreases with 
rise of temperature. At a fixed temperature the weight of gas 
dissolved by a given volume of water of other liquid is dependent 
upon the pressure of the gas and is, in general, directly propor- 
tional to the pressure. This statement is known as Henry's 
Law. The law does not apply to very soluble gases, like am- 
monia, dissolving in water — probably because chemical union 
occurs, since we know that NH 4 OH is formed in this case (91). 

127. Heat of Solution. — If we shake some potassium nitrate 
or ammonium chloride, or indeed any one of many salts, with 
water, we find that as the substance dissolves the solution becomes 
appreciably colder. This indicates that heat is required to change 
the solid into the dissolved state. This phenomenon is analogous 
to that met with when a solid, like ice, melts. It requires 79 
calories to melt 1 g. of ice, while 115 calories are absorbed when 
1 g. of potassium nitrate dissolves. That is, we must supply 
115 calories to 1 g. of the salt, and sufficient water, in order to 
prevent a fall of temperature when solution takes place. The 
heat so required is called heat of solution. 

When any substance whatever melts, heat is required, or is 
absorbed, and we might expect, similarly, that heat will always 
be absorbed when a substance dissolves; but this is not the case. 
Many substances, upon dissolving, give out heat. In the case of a 



Water and Solutions 77 

few substances the absorption or evolution of heat upon dissolv- 
ing is very small. Common salt dissolves in water with very 
small heat absorption. 

128. Boiling-Point of Solutions. — It is very easy to show that 
a solution of a solid substance, like salt or sugar in water, boils 
at a higher temperature than pure water. This is an invariable 
rule for solutions of substances which are not readily volatile 
at the boiling-point of water. Now, we have in the first part of 
this chapter (112) considered the relationship between boiling- 
points and vapor pressures, and it will easily be understood that 
a solution will boil at the temperature at which the pressure of 
its vapor is equal to the atmospheric pressure. 

129. The Lowering of the Vapor Pressure by Dissolved Sub- 
stances. — If a solution must be heated above ioo° to raise its 
vapor pressure to that which water has at ioo°, it is clear that at 
this latter temperature the solution has a lower vapor pressure than 
pure water. It is also a fact that at every lower temperature the 
vapor pressure of a solution of an involatile substance is less than 
that of the pure solvent at the same temperature. This is a 
very important universal law. The law applies to solutions 
formed from all kinds of solvents. 

130. Deliquescence. — In the case of a very soluble substance, 
like caustic soda, the vapor pressure of the saturated solution 
may be so small that it is below the partial pressure exerted by 
the vapor usually present in the air. If such a solution is 
exposed to the air, water vapor from the air will condense in it 
until the solution has become so dilute that its vapor pressure 
is just equal to the partial pressure of the water vapor in the air. 
Moreover, if such a very soluble substance is exposed to air con- 
taining moisture, water will condense on the solid, thus convert- 
ing it slowly, first into a saturated solution, and finally into a 
dilute solution. This action is called deliquescence. We say 
caustic soda is a deliquescent substance. A little thought will 
lead to the conclusion that deliquescence is the result of two con- 
current conditions; first, the possibility of the formation, by a 
substance, of a saturated solution which has a very small vapor 
pressure as compared with pure water — a condition usually 



78 Introduction to General Chemistry 

accompanying great solubility; and, secondly, the presence in the 
air of a sufficiently great water-vapor content. No substance 
is deliquescent in a perfectly dry atmosphere, while every 
soluble substance exhibits this property in air saturated with 
water vapor. Deliquescence is, therefore, not a fixed property of 
a substance. Thus common salt is usually decidedly deliquescent 
at the seashore, where the air contains much water vapor; but 
it never shows this property in a desert region. 

In several experiments we have used caustic soda or calcium 
chloride to dry air or other gases or to absorb water vapor formed 
in the burning of hydrogen (39, 50). These drying agents are 
among the most deliquescent substances known. 

131. Efflorescence. — In paragraph 94 the peculiar behavior 
of sodium sulfate decahydrate, Na 2 S0 4 , ioH 2 0, when exposed 
to the open air was described. We are now in a position to 
understand more about this spontaneous loss of water. If a 
crystal of the hydrate is floated on the surface of mercury in a 
vaccum tube like one of those shown in figure 25, the mercury 
level is depressed more than can be accounted for by the weight 
of the crystal. Apparently the latter is giving off water vapor 
and attempting to establish a saturation pressure. This pres- 
sure is called the vapor pressure of the hydrate. As a matter of 
fact all hydrates show this same behavior, with the difference 
that each has its own characteristic vapor pressure at a given 
temperature. With increased temperature the vapor pressure 
rises. If a hydrate is exposed to air in which the partial pressure 
of water vapor is less than the vapor pressure of this substance, 
the latter will give off water to the air just as a water surface 
does to air in which the partial pressure of water vapor is below 
the saturation value for water. Along with the loss of water, 
the crystals of the decomposing hydrate crumble to a powder. 
This process is called efflorescence. It is obvious that whether 
or not a given hydrate effloresces depends not only upon its own 
vapor pressure but upon the moisture content of the air surround- 
ing it. 

132. Effect of Temperature on Solubility. — The solubility of 
a substance, that is, the amount of the substance which dissolves 



Water and Solutions 



79 



(to form a saturated solution) in a given amount of water, is 
dependent upon the temperature. Most substances are more 
soluble at a higher than at a lower temperature; but this is not 
always the case, as the solubility of some substances decreases 
with rise of temperature. In fact, gases are always less soluble 
at a higher temperature. 




+Q so eo 



Fig. 27 



The change of solubility with change of temperature can 
most easily be expressed graphically, that is, by means of 
so-called solubility curves. The accompanying diagram (Fig. 2 7) 
illustrates the method and gives the curves for water solutions of 
several substances. 

133. Effect of Crystalline Form on Solubility. — Sodium sul- 
fate has the formula Na 2 S0 4 . By the action of water we may 



80 Introduction to General Chemistry 

readily obtain the hydrate Na 2 S0 4 *ioH 2 (96), which can easily 
be recrystallized from water, as described under " Supersaturated 
Solutions" in this chapter. We see that the solubility curve for 
Na 2 S0 4 - ioH 2 rises rapidly until a temperature of 33 is reached. 
At this temperature the crystals melt and at the same time 
decompose into Na 2 S0 4 and H 2 0, thus: 

Na 2 S0 4 • ioH 2 -> Na 2 S0 4 + ioH 2 0. 

Above 33 we have the solubility curve of anhydrous Na 2 S0 4 
which is a different chemical substance from its hydrate. Thus 
we see that there are for the anhydrous salt and its hydrate two 
distinct solubility curves, and that these intersect at a point 
for which the temperature is that at which the hydrate changes 
into the anhydrous substance. This is a typical case. Each 
hydrate of a substance has its own solubility curve; but these always 
intersect at the point corresponding to the temperature at which^one 
substance changes into the other. The difference in solubility is 
due to the fact that each has its own characteristic crystalline 
form. 

134. Heat of Solution and Changes of Solubility with 
Temperature. — A question which will now very naturally occur 
to the student is: Why should the solubility of various sub- 
stances change with temperature in different ways? Although a 
complete and satisfying answer cannot be given to this question, 
it is possible to find a connection between the shape of the solu- 
bility curve of a substance and another fundamental property. 
It will be recalled that potassium nitrate absorbs much heat 
upon dissolving in water, and we notice that its solubility curve 
rises rapidly with temperature. Sodium chloride dissolves with 
but slight absorption of heat, and its curve is nearly horizontal. 
Finally, when it is known that anhydrous sodium sulfate, 
Na 2 S0 4 , dissolves at temperatures above 33 with production of 
heat, and that its curve falls with rising temperature, the general 
law becomes apparent. These are typical cases. If any sub- 
stance dissolves with absorption of heat, its solubility curve 
rises with rise of temperature. If it dissolves with evolution of 
heat, then the curve falls with rise of temperature. The frac- 



Water and Solutions 81 

tional change of solubility with rise of i° of temperature is in 
general proportional to the heat of solution. In every case that 
change of solubility which will absorb heat will take place when the 
temperature is raised. This will involve a decrease of solubility 
with rise of temperature, in the case of a substance like Na 2 S0 4 , 
above 33 , since, if heat is evolved when the substance dissolves, 
heat is absorbed in equal amount when the same weight of the 
substance crystallizes out of a solution. 

In some cases where heat is evolved when a substance is 
dissolved, the observed heat is the result of the union of the solid 
with water to form a hydrate, which may dissolve with a small 
absorption of heat. In such cases the solubility of the hydrate 
increases with rise of temperature in strict accord with the law. 
For example when anhydrous calcium chloride is dissolved in 
water the mixture gets very hot. The saturated solution 
deposits crystals of CaCl 2 , 6H 2 on cooling. This hydrate dis- 
solves in water with absorption of heat and its solubility increases 
with a rise in temperature. The heat given out on dissolving the 
anhydrous salt is the excess of the heat produced in the reaction 

CaCl 2 +6H 2 = CaCl 2 ,6H 2 

above the heat absorbed in the dissolving of the hydrate CaCl 2 , 
6H 2 0. 

135. Two Apparent Kinds of Solubility. — In cases of ordinary 
solubility, evaporation of the water leads to the recovery un- 
changed of the substance originally dissolved. In other cases, 
evaporation of the solution obtained by the apparent dissolving 
of a substance leaves an entirely different substance. For 
example, if we throw a piece of sodium on water the former soon 
disappears and a solution results (40, 88) . We might be inclined 
to say that the sodium has dissolved in the water; but there is 
another way of looking at the matter. We know that in this 
case a chemical change has occurred, as represented by the 

equation 

2 Na+ 2 H 2 -> 2 NaOH+H 2 . 

Furthermore, we know that by evaporation of the solution we 
get sodium hydroxide and not sodium; for this reason it seems 



82 Introduction to General Chemistry 

more logical to say that sodium and water react to give sodium 
hydroxide, which then dissolves in water, than to say that 
sodium itself is soluble in water. In fact, we know nothing 
about the solubility of sodium in water, since the two react as 
soon as they are brought into contact. We know a very great 
number of cases analogous to this one, and in all of them we 
recognize that we have to deal with chemical changes which 
give rise to soluble products. 

136. Normal Solutions. — In the neutralization of hydro- 
chloric acid by sodium hydroxide, which takes place according 
to the equation 

HCl+NaOH -> NaCl+H 2 0, (89) 

one formula weight of the acid (36.5 g.) requires one formula 
weight of the base (40 g.). If we make a solution of the acid of 
such concentration that 1 liter contains 36.5 g. of hydrogen 
chloride, and also make a solution of the base containing 40 g. 
of sodium hydroxide per liter, then upon mixing the liter of the 
acid solution with the liter of the basic solution exact neutraliza- 
tion will take place. It follows, of course, that, to neutralize a 
given volume of such an acid solution, exactly the same volume 
of the basic solution will be required. We call such solutions 
normal solutions. 

If we wish to make a solution of nitric "acid of such concentra- 
tion that 1 liter of it will exactly neutralize 1 liter of normal 
sodium hydroxide, we see, in accord with the equation 

HN0 3 +NaOH->NaN0 3 +H 2 0, (104) 

that one formula weight of HN0 3 must be contained in 1 liter 
of the solution. This gives a normal solution of nitric acid. 

Now the case is a little different if a normal solution of sul- 
furic acid is to be made, since in this case we have 

H 2 S0 4 + 2 NaOH -> Na 2 S0 4 + 2H 2 0. (94) 

We see that one formula weight of sulfuric acid neutralizes two 
formula weights of sodium hydroxide, so that to neutralize 
1 liter of normal sodium hydroxide, which contains but one 
formula weight of the base, only one-half a formula weight 



Water and Solutions 



83 



t5CC 

AJ»G 



Fig. 2! 



/ 



(| of 98 g. or 49 g.) of sulfuric acid is required. Therefore if we 
dissolve 49 g. of the acid in sufficient water to make a liter of 
solution, this liter of acid solution will just neutralize 1 liter of 
normal sodium hydroxide. We call the sulfuric 
acid solution so made also a normal solution. 

A normal solution of potassium hydroxide, 
KOH, would contain one formula weight (56 g.) 
per liter (106). A normal solution of any acid 
always neutralizes an exactly equal volume of a 
normal solution of any base. The term " normal" 
is usually abbreviated N, so that for a normal 
solution by hydrochloric acid we write N HC1. 
Normal solutions are of great importance in prac- 
tical work. Suppose we wish to know the concen- 
tration of a given solution of sodium hydroxide. 
We take, with a pipette (Fig. 28), a carefully 
measured volume, say 20 c.c, add to it sufficient 
litmus solution to produce a pale blue color, and then from a 
measuring tube, called a burette (Fig. 29), run in a normal solu- 
tion of hydrochloric or other acid until the color just changes 
from blue to red. A little practice enables one to find, to within 

one drop or less, the volume of acid 
required. Let us say 42 c.c. of 
N HC1 was required for the 20 c.c. 
of NaOH solution of unknown con- 
centration. Our problem is to find 
the weight of sodium hydroxide in 
the 20 c.c. of solution taken. . Now, 
42 c.c. of N acid will neutralize 
42 c.c. of N sodium hydroxide, of 
which 1 liter ( = 1,000 c.c.) contains 
40 g. of sodium hydroxide. There- 
fore the weight of sodium hydroxide in the 20 c.c. taken = o . 042 X 
40 g. = 1 . 68 g. We also see that the sodium hydroxide solution 
is 42/20 = 2. 1 times as concentrated as a normal solution of this 
base. We express its concentration by saying that it is 2.1 
times normal in concentration. 




Fig. 



29 



84 Introduction to General Chemistry 

It is often convenient in practice to use solutions of J, |-, T V, 
or some other fraction of normal; we call these half -normal 

(— j, one-fifth normal (— ), and one-tenth or deci-normal ( — J, 

respectively. 

137. Acidimetry and Alkalimetry. — The analyses of acids 
and bases by means of normal solutions are called respectively 
acidimetry and alkalimetry. The act of running in a solution 
from a burette until the neutral or end-point is reached is called 
titration. The volume of solution used is called the titer. 
Instead of litmus we may use some other colored substance to 
indicate the end-point; such a substance is called an indicator. 
Other useful indicators are methyl orange, phenolphthalein, 
and Congo red. 

138. Problems. — 

N 

1. How many c.c. of — nitric acid are required to neutralize 

50 c.c. of normal potassium hydroxide? (107) 

N 

2. How many c.c. of — sodium hydroxide are required to 

N 
neutralize 20 c.c. of — sulfuric acid? (94) 

3. If 16 c.c. of a solution of sulfuric acid of unknown concen- 

tration requires for its neutralization 36 c.c. of — potassium 

hydroxide, (a) what is the weight of sulfuric acid in the 16 c.c. 
taken? (b) what is the weight of sulfuric acid in 1 liter of this 
acid? (107) 

139. The Formation of Water. — We have, in earlier chapters, 
learned various ways in which water can be formed chemically. 
We may enumerate these by way of review. 

Water is formed — 

1. By the burning of hydrogen: 

2H 2 +0 2 ->2H 2 0. 

2. By the burning of a compound of hydrogen, for example, 

methane : 

CH 4 + 2 2 -> C0 2 + 2H 2 0. (86) 



Water and Solutions 85 

3. By the oxidation of hydrogen or its compounds by means 
of combined oxygen, as, for example, when ammonia is passed 
over hot copper oxide : 

2NH3+3C11O -> 3 H 2 0+ 3 Cu+N 2 . (84) 

4. By the union of acids and bases, whereby a salt and water 
are always formed; for example: 

HCl+NaOH -> NaCl+H 2 0. (89) 

5. By the decomposition or dissociation of various unstable 
compounds, as, for example, sodium sulfate decahydrate into 
the anhydrous salt and water: 

Na 2 S0 4 • ioH 2 -> NaS0 4 -f- ioH 2 0. (96) 

140. The Chemical Reactions of Water. — We have also 
studied some of the important kinds of reactions in which water 
takes part. We may now summarize these as follows: 

1. Water unites with salts to form hydrates, thus: 

Na 2 S0 4 + ioH 2 -> Na 2 S0 4 - ioH 2 0. (96) 

2. Ammonia and water unite to form ammonium hydroxide: 

NH 3 +H 2 -> NH 4 OH. (91) 

4. Water acts upon some metals to give hydroxides and 
hydrogen. Thus, sodium and cold water react very easily, 
giving sodium hydroxide and hydrogen: 

2 Na+ 2 H 2 -> 2NaOH+H 2 . (88) 

Magnesium does not act readily on cold water, but burns vigor- 
ously in steam giving the hydroxide and hydrogen: 

Mg-f- 2 H 2 0^ Mg(OH) 2 4-H 2 . (28, 86) 



CHAPTER IX 
ACIDS, BASES, AND SALTS.— II 

141. New Acids, Bases, and Salts. — The present chapter will 
treat of three new acids, carbonic, H 2 C0 3 , acetic, C 2 H 4 2 , and 
phosphoric, H 3 P0 4 , and the bases derived from the elements 
magnesium, calcium, barium, zinc, iron, aluminum, copper, 
silver, lead, and mercury, together with the more important salts 
which these bases form with the three acids studied in the first 
chapter on acids, bases, and salts, as well as with the three acids 
above-mentioned . 

142. The Action of Water on Magnesium Oxide : Magnesium 
Hydroxide, Mg(OH) 2 . — All of the three bases studied in the first 
chapter on " Acids, Bases, and Salts" are readily soluble in water. 
We shall next consider one which dissolves in water only to a very 
slight extent. If we shake, with water, a little magnesium 
oxide (11, 80), obtained by burning magnesium, we find that the 
solution will turn red litmus blue, although but a small amount 
of the magnesium oxide has dissolved in the water, the larger 
part having remained undissolved. It has been found by careful 
experiment that magnesium oxide and water unite when brought 
together, giving a single new compound, the composition of 
which is represented by Mg0 2 H 2 , which we may also write 
Mg(0H) 2 , and call magnesium hydroxide. This is a white sub- 
stance, with which the student may already be familiar under the 
name of " milk of magnesia." It is extensively used in medicine. 
It is to be classified as a base, since, like sodium hydroxide, it 
colors litmus blue and neutralizes hydrocloric acid. The equa- 
tion for the action of water on the oxide is 

MgO+H 2 0->Mg(OH) 2 . 

143. The Action of Hydrochloric Acid on Magnesium Hy- 
droxide: Magnesium Chloride, MgCl 2 . — Magnesium hydroxide 
is but very slightly soluble in water. However, if we add hydro- 
chloric acid to the magnesium hydroxide formed from the mag- 

86 



Acids, Bases, and Salts — II 87 

nesium oxide and water until the solution just turns litmus red, 
we find that all of the solid dissolves, giving a clear, colorless 
solution which if left to evaporate in an open vessel will deposit 
colorless crystals. An investigation of this new substance shows 
that it is a compound of magnesium, chlorine, hydrogen, and 
oxygen, in the proportion indicated by the formula MgCl 2 '6H 2 0. 
This hydrate of magnesium chloride is formed as a result of the 
following two reactions: 

Mg(OH) 2 + 2HCI -> MgCl 2 + 2 H 2 0, 
MgCl 2 +6H 2 0-> MgCl 2 -6H 2 0. 

144. Magnesium Sulfate, MgS0 4 . — If we now add diluted 
sulfuric acid to some magnesium hydroxide mixed with water, 
until all of the solid has dissolved and litmus shows the solution 
to be neutral, we may obtain from the solution by careful 
evaporation crystals of magnesium sulfate having the formula 
MgS0 4 *7H 2 0, a substance much used in medicine and known as 
Epsom salts. The reaction occurs according to the equation: 

H 2 S0 4 +Mg(OH) 2 -> MgS0 4 +2H 2 0, 

the MgS0 4 then combining with water from the solution, thus: 

MgS0 4 +7H 2 -> MgS0 4 -7H 2 0, 

to form the hydrate. The latter, when heated, readily disso- 
ciates into MgS0 4 and 7H 2 0, a fact which may be expressed thus: 

MgS0 4 • 7H 2 -> MgS0 4 + 7 H 2 0. 

This, as we see, is just the reverse of the preceding reaction. 
The reactions of hydrates in solution are of course the same as 
those of the anhydrous salts, since solutions of the two cannot 
be distinguished. In what follows, the discussion of the hydrates 
formed will be omitted except in a single important instance. 

145. Magnesium Nitrate, Mg(N0 3 ) 2 . — Magnesium hydroxide 
is readily neutralized by nitric acid, with the formation of 
magnesium nitrate, which forms white crystals very easily 
soluble in water: 

Mg(0H) 2 + 2HNO3 -> Mg(N0 3 ) 2 + 2H 2 0. 



88 Introduction to General Chemistry 

Magnesium oxide and dilute hydrochloric acid react to give 
magnesium chloride, which is the same compound as that formed 
from magnesium hydroxide and the same acid. The equation 
for the reaction is 

MgO+ 2HCI -> MgCl 2 +H 2 0. 

The corresponding reactions take place with sulfuric and with 
nitric acid, and are represented by the equations 

MgO+H 2 S0 4 -> MgS0 4 +H 2 0, 
MgO+ 2HNO3 -> Mg(N0 3 ) 2 +H 2 0. 

146. Monacid and Diacid Bases: Valence. — If we compare 
the formula of magnesium chloride, MgCl 2 , with that of sodium 
chloride, NaCl, or potassium chloride, KC1, we see that, in the 
first case, one symbol weight of the metal is combined with two 
symbol weights of chlorine, while in the other two cases one 
symbol weight of metal is combined with but one symbol weight 
of chlorine. In the cases of the neutralization of the hydroxides 
by hydrochloric acid we found that one formula weight of mag- 
nesium hydroxide required two formula weights of hydrochloric 
acid (143) ; while one formula weight of the hydroxide of either 
sodium or potassium required but one of hydrochloric acid (102, 
107). For this reason we call sodium and potassium hydroxides 
monacid bases, and magnesium hydroxide a diacid base. We 
also make use of the term valence in referring to facts like those 
just mentioned, saying that the valence of sodium or potassium 
is one, while that of magnesium is two, or that sodium and 
potassium are univalent, while magnesium is a bivalent element. 
Since hydrogen chloride has the formula HC1, we say that hydro- 
gen has a valence of one, and we also say that the valence of 
chlorine is one. 

147. Radicals and Their Valence. — We have already become 
acquainted with several ammonium compounds, as, for example, 
the chloride NH 4 C1 (91) and the nitrate NH 4 N0 3 (105). We 
call the combination NH 4 the ammonium radical; it has never 
been obtained as a separate substance, but is known only as a 
component of ammonium compounds. We know of many other 
such radicals, one of which is met with in sulfuric acid and sul- 



Acids, Bases, and Salts — 77 89 

fates, where we have found that sulfur and oxygen are always 
present in the ratio represented by S0 4 . Here again we have a 
radical which is found in many salts, the sulfates, but is not 
known as a separate chemical substance. In nitric acid and 
the nitrates we have the radical N0 3 . A radical is composed of 
two or more elements united in a definite proportion; con- 
sequently the composition of a radical can always be represented 
by a formula. We may consider that the combination of nitro- 
gen and hydrogen, NH 4 , taken as a radical, has a valence of one, 
since the weight of nitrogen and hydrogen represented by NH 4 
taken once unites with one symbol weight of chlorine, giving 
NH 4 C1. Since sulfuric acid, H 2 S0 4 , forms such salts as Na 2 S0 4 , 
K 2 S0 4 and (NH 4 ) 2 S0 4 , we say that the sulfate radical, S0 4 , has 
a valence of two, a fact which is also shown by the existence of 
such salts as NaHS0 4 , etc. Now if magnesium has a valence 
of two and the sulfate radical has also the valence of two, we 
see in the fact that magnesium sulfate has the formula MgS0 4 , a 
broader meaning of the term valence. And so chemists often 
speak of the two valences of magnesium being satisfied by the 
two valences of the sulfate radical. The subject of valence will 
be considered again at the end of this chapter (183). 

148. Zinc and Its Salts. — Zinc is an element and a very 
important metal; it is known in commerce as spelter, and is used 
in enormous amounts in making galvanized iron, which is iron 
coated with metallic zinc, in making brass, whose other com- 
ponent is copper, and for many other purposes. Zinc will burn 
when strongly heated in the air or in oxygen, giving a white 
oxide, the reaction being represented by the equation 
2Zn+0 2 ->2ZnO. 

Zinc oxide is used extensively in making white paint. 

It will be recalled that magnesium burns, giving an oxide 

(11, 80), and that this oxide reacts with acids giving salts, thus: 

MgO+2HCl^MgCl 2 +H 2 0. (143, 145) 

Zinc oxide behaves like magnesium oxide when treated with 
hydrochloric acid, giving zinc chloride, thus: 
ZnO+ 2HCI -> ZnCl 2 +H 2 0. 



90 Introduction to General Chemistry 

Zinc chloride is a salt which dissolves very readily in water, 
giving a clear, colorless solution. Zinc oxide gives zinc sulfate 
and zinc nitrate as follows : 

ZnO+H 2 S0 4 -> ZnS0 4 +H 2 0, 
ZnO+2HN0 3 -> Zn(N0 3 ) 2 +H 2 0. 

These are white salts, also easily soluble in water. 

149. The Action of Hydrochloric Acid on Zinc. — If we pour 
some hydrochloric acid on zinc we observe a vigorous reaction; 
the zinc dissolves and a gas which proves to be hydrogen is given 
off. If, after the zinc has all dissolved, we evaporate the solution, 
we obtain a white solid which is found to be zinc chloride. The 
reaction is represented thus: 

Zn+ 2 HCl^ZnCl 2 +H 2 . 

Comparing this equation with that for the action of zinc oxide 
on hydrochloric acid, we see that in the latter case the hydrogen 
of the acid, instead of passing off as gas, unites with the oxygen 
of the zinc oxide, giving water. 

We might expect that metallic magnesium and hydrochloric 
acid would act thus : 

Mg+ 2 HCl->MgCl 2 +H 2 , 

and it is easy to show by experiment that this is the case. With 
dilute sulfuric acid these metals behave as follows: 

Mg+H 2 S0 4 ->MgS0 4 +H 2 , 
Zn+H 2 S0 4 -> ZnS0 4 +H 2 . 

In making hydrogen in the laboratory we usually use zinc and 
hydrochloric acid. 

150. Marble and Other Compounds of the Element Cal- 
cium. — Let us now consider the chemical behavior of marble. 
If we place some lumps of marble in a hard glass tube and heat 
strongly, a gas is given off, while the lumps change but little in 
appearance. This gas causes limewater to turn milky; it is 
carbon dioxide (19). If the lumps left after heating the marble 
are moistened with water, they grow very hot, swell up, and 
crumble to a white powder. It is evident therefore that the 



Acids, Bases, and Salts — II 91 

marble has been changed chemically by the heating. The solid 
left after the heating is the common substance, quicklime. The 
action of water upon quicklime is called slaking. If the slaked 
lime is shaken with a large amount of water, not much seems to 
dissolve; but if we filter the mixture, a clear, colorless solution 
is obtained. If some carbon dioxide gas is run into this clear 
solution it turns milky, because this solution is limewater (18), 
of which we have so often made use. If we test the limewater 
with litmus we find that it turns the latter blue, showing the 
limewater to be a solution of a base. This base reacts with acids 
to form salts. All of these products contain an element called 
calcium, whose symbol is Ca. Calcium is a brassy-looking metal, 
which will readily burn with a bright light if heated in air or 
oxygen, giving calcium oxide: 

2Ca+0 2 ->2CaO. 

Calcium oxide, CaO, is quicklime; but the latter is never 
made practically in this way, because metallic calcium is too 
expensive, and because the oxide is made very cheaply by heating 
marble or, more often, limestone, which is an impure form of the 
same compound as marble. By heating marble, CaO and C0 2 
are formed, and nothing else. By finding the percentage of 
each we can easily calculate the formula for marble to be CaC0 3 , 
which is called calcium carbonate; the effect of the heating is, 
therefore, represented thus: 

CaC0 3 ->CaO+C0 2 . 

151. Calcium Hydroxide, Ca(OH) 2 . — As has been stated, 
when water acts on calcium oxide or quicklime, we get calcium 
hydroxide or slaked lime, a solution of which is called limewater: 

CaO+H 2 0->Ca(OH) 2 . 

This reaction is analogous to the action of water on magnesium 
oxide, which was studied earlier (142). The action of hydro- 
chloric acid on calcium hydroxide gives calcium chloride and 
water: 

Ca(OH) 2 + 2 HCl -> CaCl 2 +2H 2 0. 



92 Introduction to General Chemistry 

We are now in position to understand the cause of the milki- 
ness produced when carbon dioxide acts on limewater. The 
white solid formed is really calcium carbonate, CaC0 3 . The 
equation is 

Ca(OH) 2 +C0 2 -> CaC0 3 +H 2 0. 

152. Carbonic Acid, H 2 C0 3 . — If we pass carbon dioxide into 
water, a solution results which has faint acid properties. This 
solution is in fact the well-known plain soda served at soda 
fountains. The dissolved carbon dioxide and water partially 
combine to form an acid called carbonic acid : 

C0 2 +H 2 0->H 2 C0 3 . 

Therefore we may then consider that it is this acid which 
neutralizes the base calcium hydroxide, thus: 

Ca(OH) 2 +H 2 C0 3 -> CaC0 3 +2H 2 0. 

Calcium carbonate is a salt which is almost insoluble in water. 
In fact, salts exhibit all degrees of solubility in water. Some, 
like zinc chloride, dissolve in less than their own weight of water; 
others, like common salt, are much less soluble; while many, like 
calcium carbonate, are very nearly insoluble in water. 

153. Calcium Sulfate, CaS0 4 . — Calcium hydroxide and sul- 
furic acid form calcium sulfate and water : 

Ca(OH) 2 +H 2 S0 4 ^ CaS0 4 + 2 H 2 0. 

We find by experiment that the calcium sulfate so formed dis- 
solves very slightly in water, 100 c.c. of water dissolving but 
one-fourth of a gram of the salt. On the other hand, calcium 
chloride is very soluble in water. If we add to a solution con- 
taining, say, 5 or 10 per cent of calcium chloride, a sufficient 
amount of sulfuric acid, we observe that a large amount of a 
white powder forms in the mixed solutions and soon settles to 
the bottom of the vessel, leaving a clear, colorless liquid above. 
The white powder proves to be calcium sulfate, which is formed 

thus: 

CaCl 2 +H 2 S0 4 -> CaS0 4 +2HCl. 



Acids, Bases, and Salts — 77 93 

154. Precipitation. — We often encounter chemical reactions 
in which, as in the action between calcium chloride and sulfuric 
acid, a solid is formed upon bringing together two solutions. A 
solid so thrown down is called a precipitate ; and we speak of the 
precipitation of calcium sulfate. The formation of insoluble 
calcium carbonate by the action of carbon dioxide on limewater 
is another example of precipitation. 

155- Gypsum and Plaster of Paris. — Calcium sulfate occurs 
in nature as the mineral gypsum, CaS0 4 *2H 2 0, which, if the 
water of hydration is driven off by heat, is converted into the 
well-known plaster of Paris: 

CaS0 4 - 2H 2 -> CaS0 4 + 2H 2 0. 

When powdered plaster of Paris is mixed with enough water to 
form a paste, it sets in the course of an hour into a solid mass 
which retains the form of the vessel or mold which holds it. 
Plaster casts are made in this way. The setting is due to the 
formation of interlacing crystal filaments of the hydrate 
CaS0 4 * 2H 2 0, formed by a reversal of the action by which plaster 
of Paris is formed from gypsum: 

CaS0 4 +2H 2 -> CaS0 4 - 2 H 2 0. 

156. Calcium Bicarbonate and Hard Water. — A very inter- 
esting and important reaction occurs when carbon dioxide is 
passed for a long time into a sufficiently dilute solution of calcium 
hydroxide (limewater) . At first a milkiness appears, due to the 
formation of calcium carbonate: 

Ca(OH) 2 +C0 2 -> CaC0 3 +H 2 0. 

If we continue to pass in carbon dioxide, the precipitate slowly 
dissolves, giving finally a perfectly clear solution. If this solu- 
tion is now boiled, carbon dioxide gas is given off and a white 
precipitate is formed. These facts are explained in the following 
way. Carbonic acid, H 2 C0 3 , like sulfuric acid, is a dibasic 
acid (102) and can form acid salts as well as neutral salts. Just 
as sulfuric acid yields Na 2 S0 4 and NaHS0 4 , so carbonic acid gives 
Na 2 C0 3 and NaHC0 3 , sodium carbonate and sodium acid car- 
bonate, also known as bicarbonate (baking-soda). 



94 Introduction to General Chemistry 

The calcium salts corresponding to sodium carbonate and 
bicarbonate are CaC0 3 and Ca(HC0 3 ) 2 . The difference in the 
formulae of the sodium and calcium salts is due to the fact that 
the valence of calcium is two, while that of sodium is one. Now 
when carbon dioxide, in excess, acts on calcium carbonate, 
calcium acid carbonate, called also bicarbonate, is formed, and 
this being soluble in water the precipitate goes into solution : 

C0 2 +H 2 0->H 2 C0 3 
CaC0 3 +H 2 C0 3 -> Ca(HC0 3 ) 2 . 

When the clear solution so obtained is boiled the following reac- 
tion occurs: 

Ca(HC0 3 ) 2 -> CaC0 3 +H 2 0+C0 2 . 

These reactions take place extensively in nature. Natural 
waters, e.g., those of springs and rivers, contain dissolved carbon 
dioxide, and therefore carbonic acid. Such waters passing over 
limestone, impure CaC0 3 , dissolve it and take the Ca(HC0 3 ) 2 
into solution, forming so-called hard water. When boiled, as 
in a teakettle, it gives off carbon dioxide and deposits the calcium 
carbonate. 

157. Vinegar: Acetic Acid, C 2 H 4 2 . — Acetic acid is the 

principal ingredient, other than water, in vinegar, of which it 

constitutes about 4 per cent. The formula of acetic acid is 

C 2 H 4 2 . It neutralizes sodium hydroxide according to the 

equation 

C 2 H 4 2 +NaOH = NaC 2 H 3 2 +H 2 0. 

The salt NaC 2 H 3 2 , sodium acetate, is the only sodium salt 
which can be made from this acid. Therefore the acid radical 
of acetic acid and its salts is C 2 H 3 2 and we may write the 
formula of the acid HC 2 H 3 2 to indicate that only one of the four 
hydrogen atoms of a molecule is replaceable in salt formation. 

Pure acetic acid is a colorless liquid, miscible with water in 
all proportions. It is monobasic and forms with most bases salts 
called acetates. 

158. Bone Ash: Calcium Phosphate, Ca 3 (P0 4 ) 2 . — When 
bones are burned only the gelatinous matter and connective 
tissue are removed; the white material which is left is called 



Acids, Bases, and Salts — II 95 

bone ash and consists essentially of calcium phosphate, Ca 3 (P0 4 ) 2 . 
If powdered bone ash, which is practically insoluble in water, is 
stirred with somewhat diluted sulfuric acid, the following reac- 
tion occurs : 

Ca 3 (P0 4 ) 2 +3H 2 S0 4 ->3CaS0 4 +2H 3 P0 4 . 

The calcium sulfate formed is difficultly soluble in water and may 
be filtered out, giving a clear, colorless filtrate containing dis- 
solved phosphoric acid, H 3 P0 4 . 

159. Phosphoric Acid : a Tribasic Acid. — This acid is a white 
crystalline solid, which is very soluble in water, frequently com- 
ing on the market in the form of a very concentrated solution of 
syrupy consistency. Its dilute solution has a pleasant sour taste 
and turns litmus red. With suitable proportions of sodium 
hydroxide it yields the three salts, Na 3 P0 4 , trisodium phosphate, 
Na 2 HP0 4 , disodium hydrogen phosphate, and NaH 2 P0 4 , 
sodium dihydrogen phosphate. The latter is a typical acid salt, 
having a sour taste and acid action on litmus. Phosphoric acid 
is therefore a tribasic acid. It forms with bases three series of 
salts, corresponding to those of sodium. To distinguish these 
classes of salts from one another they are called primary, second- 
ary, and tertiary, that with the smallest proportion of base being 
the primary and that in which all hydrogen is replaced being the 
tertiary. 

160. The Practical Importance of Calcium Phosphate. — 
Since calcium phosphate, Ca 3 (P0 4 ) 2 , constitutes the mineral 
matter of bones, it is, of course, a substance of very great im- 
portance. Phosphates in small amounts are also indispensable 
constituents of most plants, and it is from these, especially from 
the seeds, like wheat, oats, and corn, that men and animals get 
their needed supply. Plants, in turn, get their phosphates from 
the soil, and do not thrive on soil deficient in phosphates. Such 
infertile soil may be greatly improved by the use of fertilizers 
containing phosphates. For this purpose, bone ash is often 
employed; but since bone ash is almost insoluble in water, it is 
not directly available for plant use. In order to make it available 
it is treated with sufficient sulfuric acid to convert it into 



96 Introduction to General Chemistry 

Ca(H 2 P0 4 ) 2 , usually known as calcium superphosphate, which 
is soluble in water : 

Ca 3 (P0 4 ) 2 + 2 H 2 S0 4 -> Ca(H 2 P0 4 ) 2 + 2 CaS0 4 . 

Immense deposits of calcium phosphate occur in Florida and 
Tennessee, as phosphate rock. These deposits have doubtless 
been formed in past geological ages from the bones of marine 
animals. Phosphate rock, after treatment with sulfuric acid as 
in the case of bone ash, is used in enormous quantities as a 
fertilizer. 

161. Sodium Carbonate and Bicarbonate. — The carbonates 
of sodium which were referred to above (156) may be obtained 
by passing carbon dioxide into sodium hydroxide solution; we 
get in this way either the carbonate Na 2 C0 3 , or the bicarbonate 
NaHC0 3 , according to the proportion of carbon dioxide used. 
We may consider that the gas first unites with water to form 
carbonic acid, which then reacts with sodium hydroxide according 
to the two following equations: 

2 NaOH+H 2 C0 3 -> Na 2 C0 3 + 2 H 2 0, 
NaOH+H 2 C0 3 -> NaHC0 3 +H 2 0. 

These carbonates of sodium are manufactured in immense 
quantities, as they are very important substances. In practice 
they are not made according to the reactions given, but by more 
economical processes, which will be considered later. 

162. Potassium Carbonate and Bicarbonate. — Potassium also 
forms analogous carbonates, K 2 C0 3 and KHC0 3 ; the former, 
commonly known as potash, is contained in wood ashes, from 
which it may be dissolved by water. Upon boiling down the 
solution known popularly as lye, a residue of crude potassium 
carbonate, K 2 C0 3 , remains. This, when more strongly heated 
to burn out brown tarry matters, gives white potash, so called 
from the fact that the evaporation of the lye is carried out in an 
iron pot. This lye is extensively used in the preparation of a 
crude soft soap. A purer form of potash is used in manufacturing 
liquid soaps. Common hard soap is made from sodium car- 
bonate and fats of various kinds. 



Acids, Bases, and Salts — II 97 

163. The Action of Acids on Carbonates. — If some hydro- 
chloric acid is poured on a piece of marble (150), the liquid 
appears to boil, although the temperature does not rise notice- 
ably. It is easy to show that the apparent boiling, called 
effervescence, is due to the escape of carbon dioxide gas. The 
marble dissolves completely if sufficient acid is used, and the 
evaporated solution leaves a residue of calcium chloride (151). 
The reaction is as follows: 

CaC0 3 +2HCl^ CaCl 2 +H 2 0+C0 2 . 

Similar reactions take place between calcium carbonate and 
nitric and sulfuric acids: 

CaC0 3 +2HN0 3 -> Ca(N0 3 ) 2 +H 2 0+C0 2 , 
CaC0 3 +H 2 S0 4 ^ CaS0 4 +H 2 0+C0 2 . 

In fact, the carbonates of other elements all show this kind of a 
reaction with these acids; for example: 

NaHC0 3 +HCl-> NaCl+H 2 0+C0 2 , 
K 2 C0 3 + 2 HN0 3 -> 2 KN0 3 +H 2 0+C0 2 . 

In general, carbonates are decomposed by acids. 

164. Barium Sulfate: a Test for Sulfates. — The element 
barium resembles calcium (150) very closely in its behavior. 
Let us consider just one of its reactions at present, leaving a 
study of the others until a later time. Barium sulfate, BaS0 4 , 
is a white solid which is as insoluble in water as glass; barium 
chloride, BaCl 2 , is about as soluble as common salt. If we pour 
some sulfuric acid into a clear, colorless solution of barium 
chloride, a white precipitate of barium sulfate forms at once : 

BaCl 2 +H 2 S0 4 -> BaS0 4 +2HCl. 

We should observe the similarity of this equation to that for the 
action of sulfuric acid on calcium chloride (153). 

If we add to a solution of barium chloride a solution of sodium 
sulfate, or of magnesium sulfate, or in fact of any sulfate whatso- 
ever, a precipitate of barium sulfate is formed. For example, 
with magnesium sulfate we have 

BaCl 2 +MgS0 4 -> BaS0 4 +MgCl 2 . 



98 Introduction to General Chemistry 

By means of this reaction we can tell at once whether any solu- 
tion contains sulfuric acid or a sulfate: if no white precipitate 
is formed, sulfuric acid and sulfates are absent. We call this 
a test for sulfuric acid or sulfates. 

165. Copper and Its Compounds. — The important, familiar 
metal copper is an element. We have already learned that when 
heated in air or oxygen it unites with oxygen to form copper 
oxide (32, 33, 82), a black solid: 

2Cu+0 2 ->2CuO. 

Copper oxide reacts with the corresponding acids to form the 
chloride, nitrate, and sulfate, thus: 

CuO+ 2 HCl-> CuCl 2 +H 2 
CuO+ 2HNO3 "> Cu(N0 3 ) 2 +H 2 
CuO+H 2 S0 4 -> CuS0 4 +H 2 0. 

These salts are all easily soluble in water, giving blue solutions. 
It will be recalled that calcium oxide unites with water to 
form the hydroxide, thus: 

CaO+H 2 0->Ca(OH) 2 . 

On the other hand, if we bring copper oxide and water together, 
no union takes place. This might be taken to indicate that 
copper hydroxide, which we might expect to have the formula 
Cu(0H) 2 , cannot be formed or does not exist. This, however, 
is not the case; it is a well-known substance which is easily 
obtained in another way. If we add to a solution of copper sul- 
fate a solution of sodium hydroxide, a blue precipitate of copper 
hydroxide forms. This is a blue solid which is very nearly 
insoluble in water. Its formation takes place thus: 

CuS0 4 +2NaOH-> Cu(OH) 2 +Na 2 S0 4 . 

We may also get copper hydroxide by the interaction of solutions 
of copper chloride or nitrate with sodium hydroxide : 

CuCl 2 + 2 NaOH-> Cu(OH) 2 +2NaCl. 

If the copper hydroxide*formed in the last reaction is heated 
by boiling the mixture, the blue precipitate turns black. This 



Acids, Bases, and Salts — 77 99 

change in color is due to a change of part of the hydroxide into 
the oxide and water: 

Cu(OH) 2 ->CuO+H 2 0. 

166. The Preparation of Difficultly Soluble Hydroxides. — 

Many hydroxides of elements are nearly insoluble in water. In 
such cases, the hydroxides are formed from a solution of a salt 
of the element by adding sodium hydroxide, or potassium 
hydroxide, or in many cases ammonium hydroxide, as illustrated 
by the following equations : 

CaCl 2 + 2KOH -> Ca(OH) 2 + 2KCI, 
MgS0 4 +2NaOH-> Mg(OH) 2 +Na 2 S0 4 . 

Such difficultly soluble hydroxides separate as precipitates, which 
may be filtered out. 

167. Lead and Its Compounds. — The well-known metal lead 
is an element, which is used extensively in metallic form and also 
in the form of compounds. Lead unites with oxygen directly 
when heated in air or oxygen, giving, under suitable conditions, 
the pale yellow oxide PbO, known as litharge. This oxide, like 
those of magnesium, zinc, and copper, reacts with acids to form 
salts. Thus with nitric acid we get lead nitrate : 

PbO+ 2 HN0 3 ->Pb(N0 3 ) 2 +H 2 0. 

This salt forms large, white crystals which dissolve readily in 
water to form a colorless solution. If hydrochloric acid is added 
to a solution of lead nitrate, a white precipitate of lead chloride 
is obtained: 

Pb(N0 3 ) 2 + 2HCI -> PbCl 2 + 2HNO3. 

Lead chloride is only slightly soluble in cold water, but is much 
more soluble in hot water, from which, upon cooling, it separates 
again in white needle-shaped crystals. Lead chloride is also 
obtained from litharge and hydrochloric acid : 

PbO+ 2HCI -> PbCl 2 +H 2 0. 



ioo Introduction to General Chemistry 

Upon adding dilute sulfuric acid to a solution of lead nitrate, 
a white precipitate of lead sulfate forms : 

Pb(N0 3 ) 2 +H 2 S0 4 -> PbS0 4 + 2 HN0 3 . 

Lead sulfate is nearly insoluble in hot or cold water. 

Metallic lead is acted upon very slowly by hydrochloric acid. 
With the cold dilute acid no appreciable action takes place ; with 
boiling, concentrated acid, a very slow reaction occurs, thus : 

Pb+ 2 HCl^PbCl 2 +H 2 . 

By methods to be considered later, it is possible to prepare 
an oxide of lead containing double the proportion of oxygen 
present in PbO, namely Pb0 2 , or lead dioxide. This oxide does 
not react with dilute nitric acid. When it is heated with hydro- 
chloric acid it gives lead chloride and chlorine : 

Pb0 2 + 4 HC1 -> PbCl 2 + 2H 2 0+ Cl 2 . 

If we compare this reaction with the following, 

PbO+ 2HCI -> PbCl 2 +H 2 0, 

we see that the excess of oxygen in Pb0 2 above that in PbO oxidizes 
the hydrochloric acid, forming water and setting free chlorine. 

Lead acetate, Pb(C 2 H 3 2 ) 2 '3H 2 0, is formed by dissolving 
litharge, PbO, in acetic acid. It forms colorless prismatic crys- 
tals, which are readily soluble in water. It is a poisonous salt 
and is called sugar of lead on account of its sweetish taste. 

168. Silver and Its Compounds. — Silver is so familiar a 
metal that we need not describe its properties. It is an element 
which is most extensively used in the metallic form, but which 
forms several compounds of great practical importance. The 
metal is not readily acted upon by dilute acids, with the excep- 
tion of nitric acid, with which it undergoes a complex reaction 
represented by the equation 

3 Ag+4HN0 3 -> 3 AgN0 3 +NO+ 2H 2 0. 

We need not consider this reaction critically at this time, 
although it is well worth careful study; but note that silver 



Acids, Bases, and Salts — II 101 

nitrate is an easily soluble salt, forming a colorless solution. The 
solid salt forms large white crystals. 

169. Silver Chloride, AgCl. — The addition of hydrochloric 
acid to a solution of silver nitrate produces at once a heavy white 
precipitate of silver chloride, which is almost insoluble in water: 

AgN0 3 +HCl -> AgCl+HN0 3 . 

By adding an excess of hydrochloric acid practically all of the 
silver in a solution is precipitated. The precipitate does not 
dissolve appreciably in any of the common acids. It is, however, 
very easily soluble in ammonia solution, from which it is again 
thrown down if the solution is acidified with nitric or hydrochloric 
acid. If any solution of unknown nature gives with hydro- 
chloric acid a white precipitate which is insoluble in an excess of 
the acid, but easily soluble in ammonia, from which solution it is 
thrown down by acidifying the solution with hydrochloric acid, 
it is safe to conclude that the original solution contained a salt 
of silver. This series of reactions constitutes a test for silver in 
the form of a dissolved salt. 

170. Silver Sulfate, Ag 2 S0 4 . — This salt is formed as a white 
crystalline precipitate when sulfuric acid is added to a con- 
centrated solution of silver nitrate. It is not very soluble, 1 g. 
requiring about 200 c.c. of cold water for its solution. The same 
salt is also formed by the action of hot, concentrated sulfuric 
acid on metallic silver. 

171. Silver Phosphate, Ag 3 P0 4 . — This salt is formed as a 
yellow precipitate when sodium phosphate, Na 3 P0 4 , or some 
other soluble phosphate is added to a solution of silver nitrate: 

3 AgN0 3 -f-Na 3 P0 4 -> Ag 3 P0 4 + 3 NaN0 3 . 

This yellow precipitate is readily soluble in dilute nitric acid, 
forming a colorless solution. It also dissolves easily in aqueous 
ammonia, giving a colorless solution, from which it is again 
thrown down when the solution is exactly neutralized with nitric 
acid. 

Silver may be distinguished from lead most easily by reason 
of the solubility of lead chloride in hot water, in which silver 
chloride is insoluble. 



102 Introduction to General Chemistry 

172. Silver Oxide, Ag 2 0. — The addition of sodium hydroxide 

to a solution of silver nitrate gives a black precipitate of silver 

oxide, Ag 2 0. We might expect silver hydroxide, AgOH, to be 

formed thus: 

AgN0 3 +NaOH -> AgOH+NaN0 3 . 

Possibly this is what first happens, but, if so, the hydroxide 
formed changes at once into the oxide, 

2 AgOH->Ag 2 0+H 2 0. 

It will be recalled that copper hydroxide is decomposed at the 
temperature of boiling water into the oxide and water (165). 
In the case of silver hydroxide the change takes place at room 
temperature. 

173. Iron and Its Compounds. — The element iron is the most 
important of all metals. It unites directly with oxygen at a red 
heat, forming the oxide Fe 3 4 (81). It can also form two other 
oxides, FeO and Fe 2 3 . The oxide Fe 3 4 is magnetic and is 
called magnetic iron oxide; FeO is called ferrous oxide (from 
ferrum, iron), while Fe 2 3 is called ferric oxide. Ferrous oxide 
gives, with the corresponding acids, ferrous chloride, FeCl 2 , and 
ferrous sulfate, FeS0 4 . These salts are also formed from iron 
by the following reactions: 

Fe+ 2 HCl->FeCl 2 +H 2 , 
Fe+H 2 S0 4 ->FeS0 4 +H 2 . 

In all ferrous compounds the valence of iron is two. 

The action of hydrochloric acid on ferric oxide takes place 

thus: 

Fe 2 3 +6HCl-> 2 FeCl 3 + 3 H 2 0. 

The salt FeCl 3 is called ferric chloride. It is a dark-yellow 
substance which dissolves easily in water to form a yellow solu- 
tion. On the other hand, ferrous chloride, FeCl 2 , is pale green 
and forms a pale-green solution. We cannot get FeCl 3 from 
iron and hydrochloric acid, but we do get the salt by the action 
of chlorine on ferrous chloride, 

2 FeCl 2 -fCl 2 ->2FeCl 3 , 



Acids, Bases, and Salts — 77 103 

or on iron, 

2Fe+ 3 Cl 2 ->2FeCl 3 . 

In ferric chloride the valence of the iron is three. We also know 
ferric nitrate, Fe(N0 3 ) 3 , and ferric sulfate, Fe 2 (S0 4 ) 3 . There are, 
therefore, two series of iron salts — the ferrous, in which the valence 
of iron is two, and the ferric, in which the valence is three. 

We can obtain the two hydroxides of iron, both of which are 
nearly insoluble in water, by the action of sodium hydroxide on 
solutions of ferrous and ferric salts: 

FeCl 2 + 2 NaOH -> Fe(OH) a + 2 NaCl, 
FeCl 3 + 3 NaOH -> Fe(OH) 3 + 3 NaCl. 

Ferrous hydroxide is white if pure, but is usually obtained as a 
dirty-green precipitate; this is due to partial oxidation by the 
action of oxygen of the air, with which it readily unites. Ferric 
hydroxide is a brown precipitate. These hydroxides unite with 
acids to form salts: 

Fe(OH) 2 + 2 HCl -> FeCl 2 + 2 H 2 0, 
Fe(OH) 3 + 3 HCl-> FeCl 3 + 3 H 2 0. 

174. Aluminum and Its Compounds. — The common metal 
aluminum is an element. As is well known, the metal is not 
acted upon by air or water. It reacts easily with dilute hydro- 
chloric acid, giving aluminum chloride, A1C1 3 , and hydrogen: 

2AI+ 6HCl-> 2 A1C1 3 + 3 H 2 . 

Upon evaporation, the solution deposits white crystals of the 
compound A1C1 3 *6H 2 0. It is not possible to obtain the anhy- 
drous salt, A1C1 3 by heating these crystals, for the purpose of 
driving off water, since they decompose thus: 

2 A1C1 3 - 6H 2 -> A1 2 3 +6HC1+ 3 H 2 0. 

The anhydrous chloride is formed by the action of dry chlorine 

gas on aluminum: 

2 A1+ 3 C1 2 ->2A1C1 3 . 



104 Introduction to General Chemistry 

Aluminum chloride is easily soluble in water, forming a colorless 
solution. This solution gives with ammonia a white precipitate 
of aluminum hydroxide, Al(OH) 3 , which is insoluble in water: 

A1C1 3 + 3 NH 4 0H -> A1(0H) 3 + 3 NH 4 C1. 

When heated, the hydroxide gives the oxide and water: 

2 A1(0H) 3 ->A1 2 3 + 3 H 2 0. 

Rubies and sapphires are natural forms of aluminum oxide. 
Emery, which is a valuable abrassive, is an impure form of the 
same substance.' 

175. Various Aluminum Salts. — The hydroxide is a base 
which reacts with acids to give the corresponding salts, thus: 

Al(OH) 3 + 3 HCl-> AICI3+3HA 
Al(OH) 3 + 3 HN0 3 -> A1(N0 3 ) 3 + 3 H 2 0, 
2 Al(OH) 3 + 3 H 2 S0 4 -> A1 2 (S0 4 ) 3 +6H 2 0. 

The nitrate and sulfate are easily soluble in water, giving color- 
less solutions. The well-known substance alum is potassium, 
aluminum sulfate, KAl(S0 4 ) 2 *i2H 2 0. It is obtained in large, 
colorless crystals when a solution made from potassium sulfate 
and aluminum sulfate is allowed to evaporate. The correspond- 
ing sodium and ammonium salts are well known, and have the 
formulae NaAl(S0 4 ) 2 -i2H 2 and NH 4 Al(S0 4 ) 2 -i2H 2 0, respec- 
tively. All such compounds are known as double salts ; chemists 
are familiar with a great variety of these. Other examples of 
well-known double salts are ammonium ferrous sulfate, (NH 4 ) 2 
Fe(S0 4 ) 2 *6H 2 0, and potassium cupric chloride, K 2 CuCl 4 *2H 2 0. 

176. Acid Reaction of Aluminum Salts. — Solutions of the 
chloride, nitrate, and sulfate of aluminum, and also of alum, are 
not neutral, as we might expect, but are distinctly acid in reaction. 
They also have a sour taste. On the other hand, we find that 
moist aluminum hydroxide, if it has been carefully washed free 
from the ammonia used in precipitating it, has no action on either 
blue or red litmus. It is also tasteless. Nevertheless, we call 
the hydroxide a base, because it unites with acids to form salts. 
We say, however, that it is a weak base ; and we find in general 



Acids, Bases, and Salts — II 105 

that weak bases, of which many are known, give salts whose 
solutions are acid in reaction. This is an important matter which 
will have to be studied carefully later. 

177. Acid Properties of Aluminum Hydroxide. — If we add 
sodium hydroxide to a solution of an aluminum salt, a white 
precipitate of aluminum hydroxide is first formed, just as with 
ammonia : 

AlCl 3 + 3 NaOH-> Al(OH) 3 + 3 NaCl. 

However, upon adding an excess of sodium hydroxide, we find 
that the precipitate goes into solution. If pure aluminum hy- 
droxide is dissolved in a solution of sodium hydroxide and the 
resulting solution evaporated, crystals of sodium aluminate, 
NaA10 2 , are obtained. This substance is easily soluble in water 
and is, in reality, a salt. It thus appears that aluminum hy- 
droxide acts as an acid in this case, and we might write the equa- 
tion for the action of sodium hydroxide upon it thus: 

HA10 2 • H 2 0+NaOH -> NaA10 2 + 2H 2 0. 

We find that the solution of sodium aluminate is strongly alkaline 
toward litmus, and say, therefore, that, although aluminum 
hydroxide has some acid properties, it is a very weak acid. 

Thus we see that a substance may be both a base and an acid. 
Such a substance is said to be amphoteric. Several metallic 
hydroxides are amphoteric. Thus zinc hydroxide, Zn(0H) 2 , 
forms with hydrochloric acid, ZnCl 2 , and with sodium hydroxide, 
Na 2 Zn0 2 , sodium zincate. It is of interest to note that aluminum 
hydroxide does not react with carbonic acid, and in fact no car- 
bonate of aluminum has ever been made. Now, carbonic acid 
is a very weak acid, and aluminum hydroxide is a very weak base. 
In general, we find that very weak bases do not form salts with 
very weak acids. 

Compounds of aluminum are very abundant in the earth. 
Common clay and numerous kinds of common rocks are com- 
pounds of aluminum. 

178. Mercury and Its Compounds. — We have already learned 
something of the chemical behavior of mercury and mercuric 
oxide, HgO (13, 14, 86). The oxide, which is insoluble in water, 



106 Introduction to General Chemistry 

dissolves in dilute hydrochloric acid, giving mercuric chloride, 
HgCl 2 , and in nitric acid, giving mercuric nitrate, Hg(N0 3 ) 2 : 

HgO+ 2HCI -> HgCl 2 +H 2 0. 
HgO+ 2HNO3 ■» Hg(N0 3 ) 2 +H 2 0. 

These salts form white crystals which are soluble in water. The 
soluble salts of mercury are all extremely poisonous when taken 
internally. Mercuric chloride is familiarly known as bichloride 
of mercury or corrosive sublimate, and is extensively used as a 
powerful germicide and antiseptic. 

179. The Formation of Mercuric Salts. — The nitrate can be 
made by the action of warm, concentrated nitric acid upon 
metallic mercury: 

3 Hg+8HN0 3 -> 3 Hg(N0 3 ) 2 + 2 NO+4H 2 0. 

Hydrochloric acid does not act appreciably upon mercury; but 
the chloride can be obtained by the action of chlorine on the 
metal : 

Hg+Cl 2 ->HgCl 2 . 

It is also made by heating a mixture of mercuric sulfate and 
common salt: 

HgS0 4 +2NaCl ->HgCl 2 +Na 2 S0 4 . 

The mercuric chloride formed is readily volatile and is separated 
by sublimation; hence the old name " corrosive sublimate.' 
The process of vaporization of a solid and the condensation of its 
vapor directly to the crystalline form is called sublimation. 
The sulfate is made by strongly heating mercury with concen- 
trated sulfuric acid: 

Hg+ 2 H 2 S0 4 -> HgS0 4 + S0 2 + 2H 2 0. 

180. Mercurous Salts. — The action of cold, dilute nitric acid 
on an excess of mercury gives rise to a solution of a salt having the 
formula HgN0 3 and called mercurous nitrate. A solution of 
this salt gives with hydrochloric acid a white precipitate of 
mercurous chloride, HgCl, and with dilute sulfuric acid also a 
white precipitate of mercurous sulfate, Hg 2 S0 4 ; both of these 



Acids, Bases, and Salts — II 107 

precipitates are practically insoluble in water. Thus mercury, 
like iron, forms two series of salts: the mercurous, in which the 
element has a valence of one, or is univalent, and the mercuric, 
in which it has a valence of two, or is bivalent. 

181. The Two Oxides of Mercury. — A solution of mercuric 
nitrate gives with a solution of sodium hydroxide a yellow 
precipitate of mercuric oxide, HgO: 

Hg(N0 3 ) 2 +2NaOH-> HgO+ 2 NaN0 3 +H 2 0. 

The hydroxide of mercury, like that of silver, cannot be obtained; 
we might say that it is so unstable that it changes into the oxide 
and water as soon as it is formed; in this respect it resembles 
the corresponding compound of silver. The yellow oxide, 
formed in this way, seems to differ from the red oxide, obtained 
by heating mercury in the air or in oxygen, only in being made 
up of very much smaller particles. 

A solution of mercurous nitrate gives with sodium hydroxide 
a nearly black precipitate of mercurous oxide, Hg 2 0. 

182. Calomel. — Mercurous chloride, HgCl, which is com- 
monly called calomel, is extensively used in medicine. It is a 
remarkable but well-known fact that the usual medicinal dose 
of calomel contains many times as much mercury as does a fatal 
dose of mercuric chloride. This great difference in physio- 
logical effect is in part due to the fact that while mercuric 
chloride is easily soluble in water, mercurous chloride is nearly 
insoluble. 

183. The Valencies of Radicals. — Now that we have studied 
a considerable additional number of acids, bases, and salts, 
we may again revert to the study of valence (146, 147) since it 
furnishes a key to the easy mastery of formulae, an undertak- 
ing which is as necessary to the study of chemistry as learning 
to spell is in the mastery of a language. To write the formula 
of a chloride of a metal it is only necessary to group as many 
chlorine symbols as the metal in question has valence with the 
symbol of the latter; thus the formulae of barium chloride and 
of aluminum chloride are BaCl 2 and A1C1 3 respectfully. If we 
wish to wr te the formulae of the nitrates we group the nitrate 



io8 



Introduction to General Chemistry 



radical with the metal symbol in question according to the same 
rule, thus Ba (N0 3 ) 2 and Al (N0 3 ) 3 . To write the formulae of 
sulfates, we must again group the symbols of the radical and 
the metal so that the total valence of each satisfies that of the 
other. Since the sulfate radical is bi-valent, we may have to 
use more than one symbol weight of either the sulfate or the 
metal. Thus the formula of sodium sulfate is Na 2 S0 4 ; that of 
barium sulfate is BaS0 4 ; while that of aluminum sulfate is 
Al 2 (S0 4 ) 3 . Since (P0 4 ) is trivalent, we know at once that the 
formula of sodium phosphate must be Na 3 P0 4 , that of barium 
phosphate must be Ba 3 (P0 4 ) 2 and that of aluminum, Al (P0 4 ). 
These examples are sufficient to show how the knowledge of 
valence simplifies the writing of formulae. In Table IX the 

TABLE IX 





H 


CI 
N0 3 
OH 
C 2 H 3 2 


HCl 

NaCl 

C 2 H 3 2 

NH 4 C1 

AgCl 
HgCl 


HOH (H 2 0) 
HNO3 
NaOH 
NH 4 OH 

Ag 2 
Na,S0 4 


I. 

Univalent 


Na 

K 

NH 4 

Ag 




Hg 












II. 


Mg 

Ca 

Ba 

Zn 


S0 4 
C0 3 



MgCl 2 

CaCl 2 

BaCl 2 

ZnCl 2 

FeCl 2 

CuCl 2 

PbCl 2 

HgCl 2 


H 2 S0 4 
H 2 C0 3 
CaC0 3 
BaS0 4 


Bivalent 


Fe 




CuS0 4 




Cu 




Ca(OH) 2 




Pb 




MgO 
HgO 




Hg 












III. 


Al 


P0 4 

N 


AICI3 
FeCl 3 


H 3 P0 4 


Trivalent 


Fe 


A1P0 4 
NH 3 


IV. 




C 


CC1 4 
C0 2 


CH 4 


Tetravalent 















various elements and radicals studied are classified with respect 
to their valencies, which vary from one to four. In the column 
headed by hydrogen we have the metals together with the 
ammonium radical. In the next column, headed by chlorine, 
we have the elements and radicals that unite as a rule with 
those of the first column. The last two columns contain the 
formulae of some typical compounds. 



CHAPTER X 

THE KINETIC THEORY OF MATTER AND THE MOLECULAR 
HYPOTHESIS 

184. An Old Greek Hypothesis. — The question whether a 
portion of a given substance, say a drop of water, could be sub- 
divided to an unlimited extent, and whether the smallest particle 
so produced would still differ in no way except in size from the 
original, was one which was much debated by the Greek philos- 
ophers centuries before chemistry became a science. Anaxagoras 
(b.c. 500), who held that there was no limit to the divisibility 
of matter, was opposed by Democritus (b.c. 470), who taught 
that in the imagined process of continued subdivision minute 
particles would finally be encountered which could not be cut 
in two without destroying or completely changing the nature 
of the substance; these particles were called atoms (aro/xos, a 
body which cannot be cut in two) . This idea may be illustrated 
by the following analogy. If we take a bushel of wheat, we may 
divide it into pecks, quarts, gills, etc., and yet each measure of 
the material will be a quantity of wheat; we may go still farther, 
but we will ultimately reach the single grains, which are still 
grains of wheat; but if we cut these grains in two the resulting 
parts may no longer be called wheat, since they would no longer 
possess the most remarkable property of wheat, which is that of 
growing if planted. At present we use the term molecule (from 
Latin molecula, the diminutive of moles, a mass) to mean essen- 
tially the same as the term "atom," as used by the Greeks, and 
speak therefore of the molecular theory of matter and the molec- 
ular hypothesis. The present chapter gives an account of 
this hypothesis and aims to show how it furnishes us an explana- 
tion of many important facts. 

185. The Molecules of Water. — According to the molecular 
hypothesis, a drop of water can be subdivided, and still remain 
water, only until the single molecules are reached, and no further. 
The splitting up of the molecules of water might separate the 

109 



no Introduction to General Chemistry 

smallest particles of oxygen from those of hydrogen; but the 
result would be the decomposition of the water into its elementary 
constituents. The particles of the elements oxygen and hydro- 
gen of which the molecules of water are made up are now called 
atoms. It is supposed that all of the molecules of water are 
alike in every respect, each being made up of one atom of oxygen 
and two atoms of hydrogen. The reason for this last conclusion 
is discussed in the next chapter. In general, each molecule of a 
given pure substance is just like every other molecule of that sub- 
stance. The nature of a substance is determined by the nature 
of its molecules; and this is, in turn, determined by the number 
and kind of atoms composing the molecules. We imagine mole- 
cules to be very small, since they cannot be seen with the aid 
of a microscope of the highest power. A cubic centimeter of air 
may contain an almost inconceivable number of these tiny 
particles. 

1 86. The Molecular Hypothesis Applied to Gases. — Let us 
first consider the known facts concerning gases and try to see how 
these facts can be connected with the supposition that gases are 
made up of small particles, the molecules. In the first place, 
we know that a confined gas tends to expand and exerts a pressure 
on the walls of the vessel which holds it. If we increase the 
pressure, the volume of the gas is diminished, and by applying 
a great pressure the decrease in volume may be made very great. 
We may explain this in either of two ways: first, that the mole- 
cules, like rubber balloons, are themselves compressible; or, 
second, that the molecules, which may not be appreciably com- 
pressible, are at considerable distances from one another, but 
are brought closer together when the gas as a whole is compressed. 
Let us follow up this second idea and see to what it leads. Two 
important questions now present themselves: (i) Why are the 
molecules not in contact — that is, why should they be at con- 
siderable distances from one another? (2) Why does a gas exert 
a pressure in all directions on the walls of the vessel which con- 
tains it ? 

187. Are Molecules at Rest or in Motion? — Would it make a 
difference in the state of affairs whether the molecules were at 



Molecular Hypothesis in 

rest or in motion ? Suppose they are in rapid motion : what 
would follow? Let us recall Newton's first law of motion: 
A body at rest remains at rest, and a body in motion continues to 
move with constant velocity in a straight line, unless acted upon 
by some external unbalanced force. Now, if molecules are in 
motion, and if they behave in the same manner as other bodies, 
and if, further, they are elastic — that is, if they, like rubber or 
ivory balls, can rebound from one another or from the walls of 
the containing vessel, then they will tend to continue in motion. 
Of course, in such a case as the one imagined, the molecules of the 
gas would very frequently strike one another and also the walls of 
the containing vessel; but, being elastic, they would rebound and 
continue in motion in a new direction ; and although the velocity 
of an individual molecule might be increased or decreased as the 
result of a collision with another molecule, on the whole the 
average velocity of all the molecules would be constant. 

The various elaborations of the ideas here presented 
constitute the Kinetic Theory of Matter. This hypothesis is 
the most important corollary of the molecular hypothesis. It 
has proved enormously fruitful in explaining very diverse 
phenomena and in suggesting new lines of investigation. 

188. The Cause of Gas Pressure; Boyle's Law. — The strik- 
ing of a gas molecule against the wall of the vessel would deliver a 
little blow; and if millions of molecules struck each square centi- 
meter every second the effect would be to tend to push back the 
surface. But what is this but the exertion of pressure ? If the 
molecules strike often enough, regularly enough, and close enough 
together, this pressure would seem constant and uniform. Now, 
suppose the gas to be compressed until it occupies half its original 
volume. In each cubic centimeter there would now be double 
the original number of molecules, and on each square centimeter 
of the wall of the container twice as many molecules would 
strike per second as before; so that, as the mass of each mole- 
cule and also its velocity have remained unchanged, we should 
expect just double the pressure per square centimeter; and, in 
fact, this is just what we find by experiment. Thus we see that 
by imagining a gas to be made of numerous small, rapidly moving, 



ii2 Introduction to General Chemistry 

elastic particles, the molecules, we get an explanation of gas 
pressure and of Boyle's law. We also see how it would be pos- 
sible for the molecules to be at considerable distances (com- 
pared with their own diameters) from one another without 
tending to fall together into a mass in which the molecules would 
all be permanently in contact. 

189. The Effect of Temperature on Molecular Velocity. — We 
may next consider why it is that the molecules are in motion and 
whether the average velocity of the molecules of a given gas 
can ever be changed. We know, of course, that, for a constant 
volume, the pressure exerted by a gas increases with rise of 
temperature, so that if we are to explain gas pressure as due to the 
momenta (mass X velocity) of the molecules which strike the walls 
of the container, we must suppose that an increase of tempera- 
ture increases either the mass of a molecule or its velocity, or 
both. Naturally, the number of collisions with the wall could 
not be increased unless the velocity increased. Now, it would 
seem more reasonable to think of rise of temperature as causing 
an increase in velocity than to think of it as causing an increase 
in mass; so we have only to imagine that a rise in temperature 
causes the average velocity of the molecules to increase in order to 
get a simple and satisfying explanation of the effect of tempera- 
ture on the pressure of a gas. 

On the other hand, a decrease in temperature is accom- 
panied by a decrease in pressure; and, indeed, the pressure is at 
all times proportional to the absolute temperature. This would 
imply that the pressure would be zero at the absolute zero of 
temperature. But zero pressure could only result if the mole- 
cules were completely at rest. We may suppose, therefore, that 
at absolute zero there is no molecular motion. A body when hot 
differs from the same body when cold only by reason of the more 
rapid motion of its molecules. In short, according to this way 
of looking at the matter, heat is merely the outward manifestation 
of molecular motion. 

190. The Mixing of Gases.— We have already learned 
(122-124) that, as a rule, liquids and solids do not form perfect 
mixtures (solutions) in all proportions ; that is to say, a solid or a 



Molecular Hypothesis 113 

liquid will dissolve in a second liquid only to a limited extent. 
Not so with gases : every gas will form with any proportion of 
any other gas a perfectly uniform mixture. Air, for example, is 
a perfectly homogeneous mixture of several gases. Of course, if 
some of the gases which are brought together react chemically, 
new liquid or solid compounds might be formed which would 
separate from the gaseous mixture. But for gases that do 
not react chemically we find that all gases mix perfectly in all' 
proportions. 

191. The Diffusion of Gases. — If we bring two gases into the 
same vessel without attempting to mix them we find, after a time, 
that a perfectly uniform mixture is present in the vessel. We also 
know that if a gas like ammonia is liberated at one place in a 
closed room its odor is soon perceptible everywhere in the room. 
The process of the spontaneous mixing of gases is called diffusion, 
and we say that the ammonia has diffused through the air of the 
whole room. It is now easy to understand how this diffusion 
takes place. The molecules of ammonia are moving in all 
directions with high velocities ; the same is true also of the mole- 
cules of the air; their complete and uniform intermingling is 
therefore inevitable. 

192. The Law of Partial Gas Pressures. — It is easily found by 
experiment that, if portions of various gases are brought together 
in the same vessel, the total pressure exerted by the gas mixture is 
the sum of the pressures that would be exerted, at the same tempera- 
ture, by the same portions of these gases if each occupied the space 
alone. This is known as Dalton's Law of Partial Pressures. 
It is not difficult to explain this law, since, in a mixture of gases, 
the molecules of a given sort will strike a given area of the wall 
just as often in the presence of other unlike molecules as in their 
absence. Each kind of molecule will therefore produce the same 
partial pressure as if the others were absent. 

193. Avogadro's Hypothesis. — The student must already 
have been impressed by the fact that all gases show great similarity 
in physical behavior. They all conform to the laws of Boyle and 
Charles. This fact, together with others which we shall consider 
later, led Avogadro, then professor of physics in Turin, Italy, to 



ii4 Introduction to General Chemistry 

suggest in 1811 that it is probable that all gases contain the same 
number of molecules per cubic centimeter. Although this sugges- 
tion received some support for the first twenty years after its 
proposal, it was then nearly forgotten until about i860, since 
which time its importance and probability have been impressed 
more and more deeply on the minds of physicists and chemists, 
so that during the last fifty years it has become one of the most 
fundamental principles of chemistry. It may be stated concisely 
thus : Equal volumes of every gas or vapor at the same temperature 
and pressure contain the same number of molecules. 

Further reasons for accepting Avogadro's hypothesis will 
be given in the next chapter. Indeed, the evidence from so 
many independent sources for the truth of this view is now so 
convincing that the hypothesis is looked upon by many as a 
statement of fact, and in consequence is referred to as Avogadro's 
Law. 

194. Gas Statistics. — Within the last few years methods have 
been found by means of which the number of molecules in 1 c.c. 
of a gas has been found with a high degree of probability and 
accuracy. At o° and 76 cm. 1 c.c. of any gas has been found to 
contain 2 . 7 X 10 19 molecules, with a probable error of less than 1 
per cent. The number of molecules in 22.4 liters of a gas under 
standard conditions is therefore 22,400 X 2 . 7 X 10 19 = 6 . 06 X 10 23 . 
The number of molecules in 1 c.c, twenty-seven millions of mil- 
lions of millions, is so immense that it is difficult for the mind to 
get any tangible conception of its magnitude. However, if we 
think of the molecules in 1 c.c. as at rest for the moment, and 
uniformly distributed in rows and layers, we should then have 
in each row of 1 cm. length v 27Xio i8 = 3Xio 6 , or three million 
molecules, a number which, although large, is at least compre- 
hensible. Then in each layer there would be three million of 
these rows, and, in the whole cubic centimeter, three million such 
layers. 

195. A Cubic Mile of Sand. — Another mental picture of the 
case may be got if we imagine 1 c.c. of gas to have been expanded 
until it occupied a cubic mile. Then each row of molecules 
would be a mile long and would contain three million molecules, 



Molecular Hypothesis 115 

spaced about y V of an inch apart. Now, a grain of fine sand is 
about yV °f an mcn m diameter; and three million such grains 
placed side by side would extend one mile. Therefore, a cubic 
mile of such sand would contain 3 X io 6 cubed or 27 X 10 18 grains, 
which is the number of molecules in 1 c.c. of gas at standard 
pressure and temperature. Since the number of molecules per 
cubic centimeter has been determined by several independent 
methods which give closely agreeing results, we may safely 
accept the value given above as being correct within 1 per cent. 

196. Some Further Conclusions. — It may now be of some 
interest to note a few additional conclusions that have been 
reached in the study of gases. Let us illustrate by means of the 
gas oxygen. We know the weight of 1 c.c. of oxygen and the 
number of molecules in 1 c.c. at standard conditions; dividing 
the first by the second gives the weight of a single molecule of 
this gas; this comes out 5-3Xio~ 23 gram. The size of a mole- 
cule has also been approximately determined and in the case of 
oxygen it turns out that the diameter of a molecule is approxi- 
mately 2 . 5 X io -8 cm. We have already seen that the average 

distance between molecules is about of a cm. = 3 . 3 X 

3,000,000 

io -7 cm. We may now ask: How far, on the average, will a 
molecule travel in a straight line before it strikes another mole- 
cule? This result can be calculated when the diameter and 
average distance apart of the molecules are known, and is called 
the free path; for oxygen it is 1 . 3 X io -5 cm. Thus we see that, 
on the average, a molecule, after one collision will travel about 
40 times (i.3Xio~ 5 -r-3.3Xio~ 7 = 4o) the average distance 
between two neighboring molecules before striking a second 
molecule. This is not surprising when we note that the average 
distance between molecules is about 13 times their diameters 
(3.3X io~ 7 -^ 2 . 5 X io" 8 = 13). 

197. The Velocity of Molecular Motion. — The average ve- 
locity with which molecules travel between collisions can be calcu- 
lated with a high degree of certainty. The velocity varies with 
the mass of the molecule and its temperature but is independent 
of the pressure. Molecules of equal masses have equal velocities 



n6 Introduction to General Chemistry 

at the same temperatures, while for those with different masses 
the velocities are inversely proportional to the square root of the 
mass. At o° the velocity of the oxygen molecules is io 4 cm. per 
second, or about 15 miles per minute. But since the free path 
of a molecule of oxygen is only 1 .3 X io~ 5 cm., it will experience 
many thousands of collisions in progressing 1 cm. At each 
collision its direction of travel will change so that its actual 
progress from a given position is far slower than its high velocity 
would indicate if no collisions occurred. 

198. The Liquid State. — As a gas is compressed at constant 
temperature its molecules are brought closer together, but other- 
wise conditions remain nearly unchanged. The mass, diameter, 
and velocity of each molecule will not be altered; only the 
average distances between the molecules and their free paths 
will be shortened. It seems probable, in fact, that the average 
kinetic energy of a molecule, which is equal to one-half the 
product of its mass and the square of its velocity (J mv 2 ), remains 
unchanged, however much the gas is compressed. If we accept 
this view, we may easily extend it to cover the liquid state, in 
which we may imagine that the molecules have the same veloci- 
ties and therefore the same kinetic energies as the molecules 
of the vapor of the liquid have at the same temperature, but that 
the crowding of the molecules is so great that their free paths 
are short compared with their diameters. However, we may 
think of the molecules as able to progress slowly from one place 
to another, although the motion will be very irregular, like that of 
persons moving about in a dense crowd. 

199. Vaporization of a Liquid. — It has already been stated 
that all of the molecules of a given gas cannot have equal veloci- 
ties; nor can a given molecule always have the same velocity, 
since at every one of the frequent collisions the velocity will be 
changed. It is only the average velocity of all the molecules 
that remains unchanged as long as the temperature remains 
constant. The velocities of the molecules of a liquid also are 
not all the same at a given instant; some will be moving much 
slower, others much faster, than the average. If a fast-moving 
molecule approaches the free surface of the liquid, it may escape 



Molecular Hypothesis 117 

into the space above the liquid, whereas a slow-moving molecule, 
under the same conditions, might not be able to escape. Now 
the passage of molecules from the liquid to the space above it is 
nothing but the evaporation of the liquid. Moreover, we see 
that the rate of escape of the molecules, and therefore the rate 
of evaporation, will be greater in proportion as the average 
velocity of the molecules is increased. Since molecular velocity 
increases with rise of temperature, we get in this way a simple 
explanation as to why heating a liquid hastens its evaporation. 
When the evaporation of a liquid goes on with a poor supply of 
heat, as when water evaporates in an open vessel, the liquid 
becomes cooler. Obviously this is due to the lowering of the 
average velocity of the molecules of the liquid because of the 
escape of the faster-moving ones. 

200. Vapor Pressure. — If a liquid is placed in a closed vessel 
which it does not completely fill, it will evaporate only until the 
pressure exerted by the vapor attains a certain value which is 
definitely determined by the temperature. For example, at 
20 the vapor pressure of water is equal to that exerted by 
17.4 cm. of mercury; at 25 it equals 23 . 6 cm. Does the water 
cease to pass into vapor when these pressures are reached ? If 
so, does this mean that molecules of water no longer pass from 
the liquid to the space above it? This would seem strange. 
Let us look at the question from another point of view. Suppose 
we have a vessel full of steam and allow it to cool. We know 
that most of the steam will condense ; only a little will remain as 
vapor. If we try to picture how this occurs, we must think of 
some of the molecules of vapor, that is, gaseous water, coming 
together first to form liquid droplets; these fall to the bottom 
and soon form a layer of liquid; other molecules then strike 
this liquid and remain as a part of it. Finally, when the tempera- 
ture of the room, say 20 , has been reached, the pressure within 
the vessel will have fallen to 17.4 cm. of mercury, and most of 
the water, but not all, will have condensed to the liquid state. 
It is important to note that at a given temperature, say 20 , the 
same final vapor pressure is reached whether steam condenses 
or water evaporates. 



n8 Introduction to General Chemistry 

201. Equilibrium between Liquid and Vapor. — A very impor- 
tant question now confronts us: Do water molecules cease to 
pass from the vapor into the liquid when at 20 the pressure 
reaches 17.4 cm.? If so, Why? Would it not seem more 
reasonable to suppose that for every molecule that passes from the 
vapor into the liquid there is another that leaves the liquid and passes 
into the vapor ? This supposed state of affairs would correspond 
to that in which the number of customers in a large shop remains 
substantially constant during a given hour of the day, by reason 
of the fact that in each minute as many persons enter the shop 
as leave it. When at constant temperature the vapor pressure 
of a liquid has reached a constant value, we say there is equi- 
librium between liquid and vapor; and it would seem from the 
discussion above that this condition does not represent a state of 
rest or inaction, but one in which two opposing actions exactly 
counteract one another. 

202. Molecular Attraction. — If we think of the matter criti- 
cally, we may wonder why molecules of cooling water vapor col- 
lect into drops. Perhaps there is a sort of attraction between the 
molecules that holds them together. If so, why should it seem 
to be more effective at lower than at higher temperatures? If, 
in reality, one molecule has some attraction for another, must 
we suppose that this attraction increases with fall of tempera- 
ture ? Would it not be sufficient to assume a constant attrac- 
tion of each molecule of water for every other ? Suppose, now, 
the vapor of water is very hot; then the molecules will be moving 
with such great velocities that if two of them collide they will 
rebound, exactly as a rubber ball, thrown downward, will 
rebound on striking the floor, although gravitational attraction 
tends to keep it on the floor. But suppose the vapor to be 
cooled; its molecules will then have smaller velocities and some 
may be moving so slowly that upon collision they remain in 
contact. Other slow-moving molecules, striking by accident 
a pair of molecules so formed, may add themselves to it, and 
in this way the droplet of water could be formed. There 
are also other reasons for assuming that molecules attract one 
another. 



Molecular Hypothesis 119 

203. The Solid State. — The most striking physical difference 
between a solid and a liquid is the rigidity of the former. This 
property of solids can most easily be accounted for by assuming 
that the molecules are not free to move about as in the case of a 
liquid, where the freedom of motion is comparable to that of 
people in a crowd, but that each molecule remains in its place 
with respect to the whole solid, as well as to its neighboring mole- 
cules. It is not necessary to think of the molecules as being 
absolutely at rest. It is more likely that each molecule has a 
vibrating motion at all temperatures above absolute zero and that, 
in fact, its kinetic energy is as great as it would be if the mole- 
cule were in the vapor state at the same temperature. 

204. Crystals. — Pure chemical substances in the solid state 
usually form crystals. The crystals of a given substance all have 
the same general form. Thus, for example, the crystals of 
common salt, when perfect, are all cubical in form, while those of 
quartz occur as hexagonal prisms. If we think of a crystal as 
built up of molecules, it is natural to wonder whether the mole- 
cules are present in haphazard fashion, like potatoes in a barrel, or 
if they may not perhaps be arranged in some systematic manner, 
like bricks in a wall or balls in a regular pile. The probability 
that the molecules of a crystal are arranged in a definite and regular 
manner is greatly increased when it is known that there are exactly 
as many types of crystalline form as there are possible regular 
arrangements of points in space. 

Within the last few years it has become possible by means 
of photographic studies made by the use of X-rays to obtain very 
precise information regarding the arrangement of molecules 
forming a crystal. As a result we now know quite definitely the 
molecular structure of a number of crystals. 

205. The Melting of Crystals.— Pure crystalline substances 
have definite melting temperatures; thus, ice melts at o°, and 
potassium nitrate at 339 . Increase of temperature must 
increase the intensity of molecular vibration; at some tempera- 
ture (the melting-point) this vibration seems to become so great 
that the systematic structure of the crystal is wrecked, leaving 
only an irregularly mixed mass of molecules, forming the resulting 



120 Introduction to General Chemistry 

liquid. Crystals cannot be heated above their melting-points. 
Ice, for example, although it may be melting on the surface, is 
never hotter than zero. 

206. Supercooling. — On the other hand, water may be cooled 
2 or 3 degrees below zero without freezing, if it is kept quiet and 
is not in contact with ice. Such supercooled water immediately 
begins to freeze if touched with a piece of ice. This phenomenon 
is a common one and is easily explained. In order that the 
formation of a crystal can start, a certain minimum number 
of molecules must come together in the proper positions. But 
this exact arrangement of the several molecules necessary may 
not readily occur, especially as immediately above the freezing- 
point (which is the same as the melting-temperature) .the mole- 
cules are vibrating so fast that they are just able to shake apart 
this regular arrangement (that is, to melt the crystal). At a 
little lower temperature the molecular motion is less and therefore 
the conditions are more favorable for the starting of crystalliza- 
tion. However, if a crystal of the substance is present, then 
supercooling does not occur, but the liquid at once begins to 
crystallize (freeze) at the temperature of its melting-point. The 
reason is obvious: now each molecule that touches the crystal 
can find its proper lodging-place, and so crystalline growth can 
continue. 

207. Solutions. — In a solution the molecules of the dissolved 
substance must be very uniformly distributed; it would seem, 
therefore, that they may be moving about freely among the 
molecules of the solvent, being carried from place to place by 
their own motions. The process of dissolving of a substance 
would closely resemble that of evaporation, and the crystalliza- 
tion of a solid from its solution would correspond to the conden- 
sation of a vapor to a liquid. In fact, we may imagine that in the 
case of a saturated solution in contact with the crystals of a 
substance we have a state of equilibrium as a result of the passage 
of molecules into and out of the solution at exactly equal rates. 



CHAPTER XI 

THE ATOMIC HYPOTHESIS AND ATOMIC WEIGHTS 

208. Dalton' s Atomic Hypothesis. — The application of the 
Atomic-Molecular Hypothesis to the explanation of chemical 
phenomena was first made by John Dal ton of Manchester in 
1803. Long before this time Bernoulli had proposed the Kinetic- 
Molecular Hypothesis as an explanation of the physical behavior 
of gases, and Daiton, knowing this view of the nature of matter, 
sought to explain the difference in solubility in water of different 
gases as due to a possible difference in size of their molecules. 
But how could this imagined difference be discovered ? At this 
date, 1803, the theory of the indestructibility of matter and the 
doctrine of elements were well established, owing to the work of 
Lavoisier, a quarter of a century earlier, as well as the labors of 
many able chemists of the intervening period. It was generally 
accepted that the formation of a substance was due to the union 
of the elements composing it and in many cases the proportions 
of the elements in a compound were already known — not very 
accurately, it is true, but at least approximately. Dalton wished 
to discover the relative weights of the ultimate particles of gases ; 
but in order to do this he would have to know, in the case of 
hydrogen and oxygen, for example, in addition to knowing the 
weight of oxygen that would combine with a given weight of 
hydrogen, the relative numbers of ultimate particles of the two 
gases that combine with one another in the formation of water. 
As Dalton had no experimental means of discovering the informa- 
tion he lacked, he simply assumed that one ultimate particle 
of hydrogen united with one ultimate particle of oxygen to give 
one ultimate particle of water, meaning by the expression ulti- 
mate particle essentially the same as the Greeks and later philos- 
ophers meant by the terms "atom" or "molecule," that is, the 
smallest possible particle of the substance. 

209. Finding the Relative Weights of Atoms. — Now, since 1 g. 
of hydrogen unites with 8 g. of oxygen to form 9 g. of water, 



122 Introduction to General Chemistry 

the ultimate particle or atom of oxygen must weigh eight times 
as much as the ultimate particle or atom of hydrogen; and the 
ultimate particle of water, in this case the molecule, must weigh 
nine times as much as an atom of hydrogen. In making such 
suppositions Dal ton also assumed that all of the atoms of hydro- 
gen were exactly alike in size, weight, and all other properties; 
that each atom of oxygen was exactly like every other atom of 
this element, but entirely different from an atom of any other 
element. Dalton knew that the same pair of elements often 
form two or more compounds in which the constituents are 
present in different proportions. This forced him to assume 
also that in such cases the atoms unite, not only one to one, but 
also one to two, or one to three, etc. In order that the student 
may have a perfectly clear notion of the matter, we may sum- 
marize by stating that Dalton assumed that a molecule of water 
is composed of one atom of hydrogen and one atom of oxygen, 
and then reached the conclusion that an atom of oxygen was 
eight times, and a molecule of water nine times, as heavy as an 
atom of hydrogen. But Dalton did not know, as we can .see 
clearly, whether one atom of hydrogen unites with one atom of 
oxygen or with two or three of oxygen, or whether two or perhaps 
three atoms of hydrogen unite with one of oxygen to form a 
molecule of water: it was all a guess. But it must also be clear 
that if we could discover the numbers of atoms of hydrogen and 
oxygen in a molecule of water we could find the relative weights of 
the two atoms, knowing the percentages of hydrogen and oxygen 
in water. Now the question is: How can we discover the num- 
ber of atoms of each kind in a molecule of a substance ? 

210. The Application of Avogadro's Hypothesis. — Suppose 
we accept Avogadro's suggestion that equal volumes of all gases 
at the same temperature and pressure contain the same number 
of molecules, and see to what conclusion we are led. Let us 
represent by N the number of molecules in 22.4 liters of any 
gas under standard conditions. Now according to Dalton's 
suggestion one molecule of a given substance will contain one, 
two, three, or some small whole number of atoms of a given 
element, but cannot, by reason of the assumed indivisible nature 



Atomic Hypothesis and Atomic Weights 123 

of an atom, contain a fraction of an atom. Let us consider the 
gas ammonia as an example. Ammonia is composed of 17.8 
per cent of hydrogen and 82.2 per cent of nitrogen, and nothing 
else. One molecule of ammonia, according to Dalton's sug- 
gestion, contains one, two, three, or four, or at least some small 
number of atoms of hydrogen. Now, if 22.4 liters of ammonia 
gas under standard conditions contain N molecules, then this 
volume of the gas must contain 1XA 7 , 2XN, 3XA 7 , or some 
small number of times N atoms. The least number of hydrogen 
atoms that could possibly be contained in 22.4 liters of ammonia 
is N, but the true number may be 2 X A 7 , which we may write 2N, 
or it may be greater, as 3 A 7 or 4N, so far as we know; only it 
must be N or some small whole number of times N if we assume 
that there are N molecules in 22.4 liters of the gas and also 
assume, with Dalton, that each molecule of the gas contains 
one, two, three, or some small number of atoms of hydrogen. In 
the case of any other gaseous compound of hydrogen we should 
conclude, according to Avogadro, that 22.4 liters of the gas 
contained N molecules and that each molecule contained one, 
two, three, or some other small number of hydrogen atoms, the 
smallest possible number being one atom of hydrogen to the 
molecule, and therefore that 22.4 liters of the gas would contain 
A 7 , 2 A 7 , or 3 A 7 , etc., atoms of combined hydrogen. 

2.1 1. The Number of Atoms and Weight of Hydrogen in 22.4 
Liters. — According to Dalton all hydrogen atoms are alike and 
each has a definite weight, so that the weight of N atoms of 
hydrogen would be a perfectly definite weight of this element. 
The weight of 2 A 7 atoms of hydrogen would, of course, be twice 
that of N atoms, etc. It seems reasonable to think that it would 
be likely to happen that in some of the gaseous compounds of 
hydrogen the molecules would contain but one atom of hydrogen 
each. In such a case 22.4 liters would contain A 7 atoms of 
combined hydrogen, having a definite weight. Now as such 
gases contain the minimum possible number of atoms of hydro- 
gen in each molecule, namely, one, and as we assume that all 
gases contain A 7 molecules in 22.4 liters, then such gases would 
contain the minimum possible weight of hydrogen in this volume. 



124 Introduction to General Chemistry 

As a matter of fact we actually find that in 22 . 4 liters of the vari- 
ous gaseous compounds of hydrogen the weight of this element is 
in no case less than 1 g. In this volume of any definite gas there is 
either no combined hydrogen or there is at least 1 g.: the minimum 
weight of hydrogen is 1 g. 

212. The Explanation of the Laws of Minimum and Multiple 
Weights. — In other gaseous compounds of hydrogen we find in 
22.4 liters larger weights of combined hydrogen, but these 
weights are then either 2 g., 3 g., or some whole multiple of the 
minimum weight. Is it not logical then to think that in such 
gases as hydrogen chloride, where the minimum weight, 1 g., 
of combined hydrogen is contained in 22 . 4 liters of the compound 
gas, the molecule contains but one atom of hydrogen, and that 
in acetylene, where 2 g. of hydrogen are found in 22.4 liters 
each molecule contains two atoms of hydrogen, while in methane 
with 4 g. of hydrogen in the same volume, there are four atoms 
of hydrogen per molecule? Undoubtedly so. We see then 
that we have in the assumptions made by Avogadro and Dalton 
the basis of an explanation of the remarkable Laws of Minimum 
and Multiple Weights, which have been discovered by experi- 
ment; and, because of the agreement between theory and fact, 
we are inclined to think that perhaps the views of Avogadro and 
Dalton are correct. In any case we cannot fail to see that these 
hypotheses are useful, and that, indeed, is the criterion by which 
the worth of any hypothesis should be judged. 

213. Application of the Explanation to Other Elements. — The 
question now arises whether the simple explanations of the 
laws of minimum and multiple weights may be applied to ele- 
ments other than hydrogen, and a very little thought will show 
that this must be the case. The minimum number of atoms 
of any given element in 22.4 liters of any of its gaseous com- 
pounds is again N, the number found in the case of those gases 
the molecules of which contain but one atom each of the given 
element. The minimum weight of this element is, of course, the 
weight of N atoms of the element. For example, we find that 
in 22.4 liters of gaseous carbon compounds the minimum weight 
of carbon is 12 g. Since this minimum weight is found in the 



Atomic Hypothesis and Atomic Weights 125 

gases carbon dioxide and methane, we conclude that but one 
atom of carbon is contained in a molecule of each. On the other 
hand, 22.4 liters of acetylene contain 24 g. of combined carbon, 
which is twice the minimum weight, from which we conclude 
that in a molecule of this gas there are two atoms of carbon. 
As we have now reached the conclusion that a molecule of methane 
contains four atoms of hydrogen and one of carbon we see that we 
have developed a method whereby we can solve the problem 
first suggested by Dalton, that of discovering the number of atoms 
of each sort in a molecule of a given substance at least in the case 
where the substance is a gas, since it is evident that the method 
used for methane is applicable to any gaseous substance. 

214. The Number of Atoms of Each Kind in a Molecule. — To 
illustrate by further examples, we may consider the cases of a 
few of the gases we have already studied. We see, by reference 
to Table IV, that in 22.4 liters of hydrogen chloride there are 
found the minimum weights of both hydrogen and chlorine and 
conclude that the molecule of this gas is made up of one atom 
each of hydrogen and chlorine. We also see by Table IV that 
the unit volume of ammonia contains the minimum weight of 
nitrogen and three times the minimum weight of hydrogen, and 
decide that in a molecule of ammonia one atom of nitrogen must 
be united with three atoms of hydrogen. In all other cases the 
reasoning is equally simple, so that the student will have no 
trouble in deciding upon the number of atoms of each kind in a 
molecule of each of the gases mentioned in Table IV. 

215. The Number and Kind of Atoms in a Molecule Shown 
by the Formula. — We are now in position to notice a most re- 
markable fact, which the following examples will illustrate. One 
molecule of hydrogen chloride contains one atom of hydrogen and 
one of chlorine, and its formula is HCl; one molecule of ammonia 
contains three atoms of hydrogen and one atom of nitrogen, and 
its formula is NH 3 ; one molecule of methane contains four atoms 
of hydrogen and one atom of carbon, and its formula is CH 4 ; 
in each case the number of atoms of each element is the same as the 
number of symbol weights of that element in the formula of the sub- 
stance! And that the same thing is true for all gases of Table IV 



126 Introduction to General Chemistry 

may readily be found by considering each separate case in the 
same way as we did those of three of the gases. In every case, 
therefore, the formula shows not only the weight of each element in 
22 . 4 liters of the gas but also the number of atoms of each element 
in one molecule of the substance. 

2 1 6. Symbol Weights and Atomic Weights. — But we may 
now inquire, Why should this be true ? To answer this question, 
we will recall that the minimum weight of any element in 22.4 
liters of its gaseous compounds is the weight of N atoms of that 
element. If N atoms of hydrogen weigh 1 g. and N atoms of 
carbon weigh 12 g., then one atom of carbon must be 12 times as 
heavy as an atom of hydrogen. In a similar way we are led to 
conclude that an atom of nitrogen is 14 times, and an atom of 
oxygen 16 times, as heavy as an atom of hydrogen. Analogous 
relations must likewise exist in the cases of all other elements; 
and, therefore, taking the weight of one atom of hydrogen as one 
or unity, the weight of an atom of any other element is repre- 
sented by exactly the same number as its symbol weight. For 
this reason a table of symbol weights is also called a table of 
Atomic Weights ; and symbol weights are usually referred to as 
atomic weights. But we must remember that the symbol 
weights may be found by simple and direct experiments, inde- 
pendently of all suppositions and hypotheses, while atomic 
weights are to be represented by the same set of numbers only 
when we assume that matter is made up of atoms which unite 
in simple ratios to form molecules of which all gases are assumed 
to contain equal numbers in equal volumes. Briefly stated, 
symbol weights are natural constants, but atomic weights are the 
probable relative weights of the atoms of which we imagine matter 
to be made up. We now may answer the question proposed in the 
first sentence of this paragraph. The number of atoms of any 
sort in a molecule is the same as the number of symbol weights of 
that element because the absolute weight of an atom of any element is 
proportional to its symbol weight. In this chapter we have seen 
how the problem which Dalton set for himself over a century 
ago is to be solved, at least as definitely as chemists know, up 
to the present time, how to solve it. The key to the solution was 



Atomic Hypothesis and Atomic Weights 127 

the hypothesis of Avogadro, which was suggested in 181 1, only 
three years after Dal ton's views first appeared in print, and 
which was rejected by Dalton himself, and was only accepted 
by the chemical world at large half a century later. 

217. Formula Weights and Molecular Weights. — It is of 
course obvious that the weight of a molecule may also be 
expressed in terms of the weight of one atom of hydrogen which 
is taken as unity. For example, if one atom each of hydrogen 
and chlorine compose a molecule of hydrogen chloride and if , as 
we have seen, an atom of chlorine weighs 35.5 times as much as 
an atom of hydrogen, then a molecule of the compound must 
weigh 36. 5 times as much as an atom of hydrogen; and we say 
therefore that the Molecular Weight of hydrogen chloride is 
36.5. The molecular weight of a gas has consequently the same 
numerical value as its formula weight, the weight of an atom of 
hydrogen being in all cases taken as unity. The conclusion 
that the relative weights of the molecules of gases are propor- 
tional to their respective formula weights follows at once from the 
assumption of Avogadro's hypothesis. But we now see also that 
the weights of gaseous molecules, briefly their molecular weights, 
are all represented by the same numbers as their formula weights 
if we choose the atomic weight of hydrogen as unity; for this 
reason it seemed logical to discuss atomic weights before molec- 
ular weights. It is also evident that the molecular weight of a 
substance must be equal to the sum of the atomic weights indicated 
by the formula. 

218. The Formulae of Some Elementary Gases. — We are now 
in position to consider the meaning of the fact that the formulae 
H 2 , 2 , N 2 , and Cl 2 were found for the four elementary gases 
studied. If we accept Avogadro's hypothesis for these as well 
as for compound gases, then the unit volume of each gas must 
also contain N molecules. But we know also that the weight 
in each case is twice the minimum weight of N atoms; for this 
reason we are forced to conclude that the unit volume of each 
gas contains 2 N atoms, and hence that each molecule contains 
two atoms. The same rule that applies to compound gases 
applies here also: the number of atoms of each element in any 



128 Introduction to General Chemistry 

molecule is the same as the number of symbol weights of that 
element. The obvious meaning of these facts is that an atom 
of hydrogen, for example, can unite with another atom of hydro- 
gen as well as with one of chlorine or some other element. But 
if this is the conclusion, you will doubtless ask: Do we have any 
other evidence of its truth? Let us see. 

219. The Union of Hydrogen and Chlorine by Volume. — It 
will be recalled that when hydrogen and chlorine gases unite to 
form gaseous hydrogen chloride, one volume of each of the 
elementary gases combines to give two volumes of the product. 
Now in two unit volumes of 22 .4 liters each of hydrogen chloride 
there must be 2 N atoms of hydrogen and 2N atoms of chlorine, 
since we cannot have less than one atom of each element in a 
molecule of the compound ; but the two unit volumes of hydrogen 
chloride are formed from one unit volume of hydrogen and one of 
chlorine each containing N molecules; and again we are led to 
the conclusion that N molecules of hydrogen or chlorine contain 
2 N atoms in each case, or that each molecule of either elemen- 
tary gas contains two atoms. 

220. Gay Lussac's Law of Combining Volumes. — The simple 
volumetric relation between hydrogen, chlorine, and hydrogen 
chloride, 1:1:2, is not an exceptional case; other gases also 
exhibit similar simple relations. Thus, two volumes of hydrogen 
and one volume of oxygen unite and, if the temperature at which 
the experiment is carried out is so high that the water remains in 
the form of steam, the latter measures two volumes; so that the 
volume relations are 2:1:2. In the burning of methane one 
volume of the gas requires two volumes of oxygen and gives 
one volume of carbon dioxide and two volumes of steam, the 
measurements all being made at a sufficiently high temperature 
in this case also to keep the steam from condensing. Or, when 
ammonia gas is decomposed, as it may be by means of electric 
sparks, two volumes of ammonia yield one volume of nitrogen 
and three volumes of hydrogen. The fact that gases and vapors 
of volatile substances always react in simple ratios by volume 
was discovered by Gay Lussac in 1808, and is known as Gay 
Lussac's Law of Combining Volumes.' 



Atomic Hypothesis and Atomic Weights 129 

221. Explanation of the Law of Combining Volumes. — The 

explanation of this law will appear if we write the equations of 
the reactions mentioned: 

H 2 +C1 2 -> 2 HC1 CH 4 + 2 2 ^C0 2 +2H 2 

I VOl. I VOl. 2 Vols. I VOL 2 VOls. I Vol. 2 Vols. 

2H 2 +0 2 ->2H 2 2NH 3 -> 3 H 2 +N 2 

2 VOls. I VOl. 2 VOls. 2 Vols. 3 Vols. I Vol. 

Again we see, as we did earlier (76, 77), that the volumes are 
the same as the coefficients of the formulae in the equations, and 
this for the fundamental reason that one formula weight always 
represents one unit volume, in the case of a gaseous or volatile 
substance. Moreover, we now understand, to cite the first 
example, that one unit volume of hydrogen containing N mole- 
cules and 2 N atoms will require 2 N atoms of chlorine or N mole- 
cules, which, according to Avogadro's hypothesis, will be found 
in one unit volume of chlorine. The reaction will then produce 
2 N molecules of hydrogen chloride, which according to the same 
hypothesis will occupy two unit volumes. Similar reasoning 
may be applied to all other cases. 

222. The Degree of Accuracy of Symbol Weights. — Before 
leaving the discussion of symbol and atomic weights we must 
consider the degree of accuracy of the statements of numerical 
results made in early chapters and summarized in Table IV. It 
is perhaps needless to point out that statements of lengths, areas, 
volumes, weights, etc., whether they refer to scientific or other 
matters, are in general more or less approximate, the degree of 
accuracy aimed at being determined by the requirements of the 
case. Thus, if a stranger in the city inquires the distance from 
the City Hall to the University and is told by a policeman that 
it is seven miles, the answer is quite as accurate as necessary. 
But such approximate statements of distance would not satisfy 
the requirements of a surveyor who wished to make an accurate 
map of the city. Up to about twenty-five years ago the most 
accurate analyses of water indicated that 2 g. of hydrogen were 
combined with 15.96 g. of oxygen. As all chemists know that 
in every analysis there is inevitably some experimental error 



130 Introduction to General Chemistry 

of greater or less magnitude, it was thought that the true weight 
of oxygen combined with exactly 2 g. of hydrogen was exactly 
16 g. It then became apparent from the new researches of a 
number of chemists that the error in the accepted results was 
greater than suspected, and, moreover, that the true proportion 
of oxygen in water was less instead of greater than the value 
found earlier, the new experiments leading to a ratio of 2 to 15 . 88, 
with a probable error of less than o. 01 g. in the weight of oxygen 
combined with exactly 2 g. of hydrogen. 

223. = i6.ooo, the Real Basis for Symbol and Atomic 
Weights. — An annoying difficulty now arose from the fact that 
far more symbol weights had been found by the analysis of oxygen 
compounds than by the analysis of compounds with hydrogen, 
owing to the greater accuracy with which the former analyses 
could be made ; so that it then became necessary for chemists to 
decide whether they should change the symbol weights of oxygen 
and all elements whose symbol weights had been found by the 
analysis of their oxygen compounds, or whether they should 
change the symbol weights of hydrogen and a few other elements. 
After much debate the former policy was adopted and the symbol 
weight of oxygen, = 16.000, kept unchanged, although this 
made it necessary to change the symbol weight of hydrogen to 
1.008. Our most accurate knowledge of the composition of 
water is expressed by the statement that 2X1. 008 g. of hydrogen 
are combined with 16.000 g. of oxygen in 18.016 g. of water, a 
fact which is also expressed by the formula H 2 0, when we consider 
that H = 1.008 g. of hydrogen and = 16.000 g. of oxygen. 
Oxygen with a symbol weight of 16.000 has thus become the real 
basis of the system of symbol and atomic weights rather than hydrogen 
with a symbol weight of unity. 

224. The Method of Finding Symbol Weights. — The symbol 
weights, and therefore also the atomic weights, of all other ele- 
ments are now based upon that of oxygen taken as 1 6 . 000 ; but 
we see by a comparison of the values given in a table of exact 
atomic weights that in no case does the exact value based on 
= i6.ooo differ greatly from the approximate value we have 
previously used. Just as more accurate analyses led to a change 



Atomic Hypothesis and Atomic Weights 131 

in the symbol weight of hydrogen, so also newer analyses have 
led and will continue in the future to lead to a more exact 
knowledge of the symbol weights of other elements. We do not 
expect, however, that the values accepted at present for the 
commoner elements will be changed by more than a few units 
in the second decimal place. Concisely stated, the matter stands 
thus: Approximate symbol weights are found in the manner 
described in chap, v, while the more exact values are fixed by the 
most painstaking analyses and syntheses, being computed on the 
basis of = i6. 000. 

225. Inexactness of the Gas Laws. — -The gas laws of Boyle, 
Charles, and Avogadro are also only closely approximate state- 
ments of the facts. For example, if the pressure on 1,000 c.c. 
of oxygen under standard conditions be exactly doubled, the 
volume will become 499.3 c.c. instead of exactly 500, as Boyle's 
law would indicate. The deviations from the simple laws are 
thought to be due to attractions between the molecules, on the 
one hand, tending to diminish the volume, and, on the other 
hand, to the fact that part of the space occupied by the gas is 
filled with the molecules themselves, so that the free space is 
reduced to less than half if the volume of the gas is, by increase 
of pressure, reduced- to half. The actual deviations from the 
simple law, PV = & constant, become negligible if gases are under 
low pressures. Then the three great laws express almost exactly 
the behavior of all gases. In other words, if the barometric 
pressure at sea-level were 0.01 of its actual value, so that our 
standard of atmospheric pressure would be o. 76 cm. of mercury 
instead of 76 cm., then we should find that not only would the 
laws of Boyle and Charles express with a high degree of accuracy 
the behavior of gases under pressures of this order of magnitude, 
but that for all gases the law of Avogadro would also hold good 
with as great a degree of accuracy as experiment would enable 
us to determine. 

226. Exactness of Avogadro' s Law for Corrected Gas Vol- 
umes. — Now, instead of trying to weigh and measure gases under 
such low pressures in attempts to study them more accurately, 
chemists have worked at ordinary pressures and then corrected 



132 Introduction to General Chemistry 

the data so obtained so as to give the results that would theoreti- 
cally have been obtained for the weights of 1 liter if the measure- 
ments had been made at very low pressures and the calculations 
made for a pressure of 76 cm. exactly according to Boyle's law. 
Working in this way, it was found that the corrected volume of 
32 g. of oxygen, the weight represented by 2 , is 22.41 liters 
at o°. It was then discovered that exactly this (corrected) 
volume of any other gas at o° contains, as nearly as the deter- 
minations could be made, just the weight of the gas which its 
formula indicates, this weight being calculated from the most 
exact symbol weights. In others, Avogadro's law would hold 
exactly at low pressures or also at ordinary pressure if the attrac- 
tions of the molecujes for each other did not exist, and if their 
own volumes were negligible as compared with the total space 
occupied by the gas. 

227. A Little Explanation and Advice. — It is not necessary nor 
desirable that the beginner in chemistry should pay much atten- 
tion to the matters discussed in the, three preceding paragraphs. 
The approximate symbol weights and the gas laws in their sim- 
plest forms are sufficiently exact for his use. It is much better 
that he should see clearly the general fundamental principles 
than that he should be perplexed and confused by the details 
and refinements that are of importance only to the specialist. 
If the beginner continues his study of chemistry he will be sure 
to encounter later these interesting topics, when he will be better 
able to appreciate and understand them ; while if he should not go 
farther than the first course, he may feel assured that he has be- 
come acquainted with the principles of most fundamental impor- 
tance. These matters are discussed here in order to explain why 
the symbol or atomic weights given in Tables of Atomic Weights, 
(see inside of back cover of this book) are not exactly the same 
as those we have used in the earlier chapters. 

228. Means of Discovering Symbol Weights. — The student 
will doubtless have received the impression from the study of the 
foregoing chapters that we can discover the approximate symbol 
weight of any element by finding the minimum weight of the 
element in the unit volume of its gaseous or vaporized com- 



Atomic Hypothesis and Atomic Weights 



*33 



pounds; and this, in fact, is true for a large number of elements 
in addition to the five included in Table IV. We shall now 
consider some facts leading to a knowledge of the symbol weights 
of a dozen elements other than the five already studied. These 
twelve elements all form volatile compounds, the densities of 
which may be determined by making experiments at sufficiently 
high temperatures and then calculating, by the laws of Boyle 
and Charles, for the standard conditions, the weight of the com- 
pound in 22.4 liters. Multiplication of the weight so found by 
the percentage of the element in question in the compound gives 
the weight of the element in 22.4 liters of the vapor, as recorded 
in Table X. 

TABLE X 



Volatile Compounds of 
Various Elements 

Antimony trichloride. . 
Arsenic trichloride. . . . 
Bismuth trichloride. . . 

Cadmium 

Chromium oxychloride 

Hydrogen iodide 

Iron carbonyl 

Lead chloride 

Mercury 

Nickel carbonyl 

Phosphorus trichloride 
Zinc chloride 



Weight of 
Elements in 
2 2. 4 Liters 



Symbol and 
Symbol Weight 



Specific Heat 



Product of 
Symbol Weight 
andSpecificHeat 



HQ-5 

75-4 

217.0 

114. o 

55-o 

127.7 

53-2 

207. 2 

202.2 

59-7 

3i-9 

63 -5 



Sb =120.2 
As = 75.0 
Bi =208.5 
Cd =112.4 
Cr = 52.1 
I. =127.9 
Fe = 55.9 
Pb =206.9 
Hg = 200 . o 
Ni = 58.7 
P = 31.0 
Zn = 65.4 



0.0503 
0.0830 
0.0303 

o-055i 
o. 1121 
0.0541 
o. 1162 
o . 0304 
o . 0308 
0.1084 
o. 2020 

0.0935 



6.0 

6.2 

6.3 

6.2 

5-8 
6.9 
6-5 
6-3 
6.1 
6.4 
6-3 
6.1 



In all cases the compounds are such as contain the mini- 
mum weight of the element the symbol of which appears in the 
table; that is to say, we do not know any other volatile com- 
pounds in the respective cases containing appreciably smaller 
weights in the unit volume. The weights so found are, therefore, 
approximately the symbol weights in each case. The exact 
symbol weight in any case is then calculated from the accurately 
determined percentage composition of some compound of the 
element with an element of exactly known symbol weight. 

229. The Product of Specific Heats and Symbol Weights. — 
There are a great many elements which do not form gaseous com- 
pounds, or compounds which are sufficiently volatile without 



134 Introduction to General Chemistry 

decomposition, to enable us to find their symbol weights in the 
manner above indicated. Very fortunately other methods have 
long been known by which the desired end can be attained. We 
shall now consider one of these methods. 

A very simple relation was discovered nearly a century 
ago, by Dulong and Petit, between symbol weights and specific 
heats of solid free elements. The amount of heat required to 
raise the temperature of a given weight of iron i° would raise 
the temperature of an equal weight of water only o. 1162 ; and 
we say, therefore, that the specific heat of iron is o. 1162. The 
specific heats of the other elements of Table X-are given in the 
fourth column. If, now, we multiply the specific heat of an 
element by its symbol weight we get the remarkable series of 
products contained in the last column of the table, where we see 
that the values are nearly the same in all cases. Does it not 
seem probable that the law which we find applying to the ele- 
ments of Table X would also hold good for other solid elements 
even though they do not form easily volatile compounds? If 
so, it is clear that in order to find the approximate symbol weight 
of an element we have only to divide 6.4 by its specific heat, which 
latter constant can in general be found by a simple, direct 
experiment. As a matter of fact, this method has been of much 
service in just this way. 

230. Interpretation of the Law of Dulong and Petit. — The law 
of Dulong and Petit is, moreover, of the greatest interest and 
importance when viewed from the theoretical standpoint. The 
product of the specific heat and symbol weight is obviously the 
quantity of heat required to raise the temperature of the symbol 
weight of an element one degree ; and this amount of heat is the 
same for one element as for another. But the symbol weights 
of various elements are the weights of equal numbers of atoms, 
and we see, therefore, that it requires equal amounts of v heat 
to raise the temperature of equal numbers of various kinds of 
atoms by one degree! The products of symbol weights and 
specific heats are generally called Atomic Heats; so that the 
Law of Dulong and Petit may be stated thus: The atomic heats 
of the solid elements are equal. 



CHAPTER XII 

THE HALOGENS AND THEIR COMPOUNDS WITH HYDROGEN 

AND METALS 

231. The Halogens. — The elements fluorine, chlorine, bro- 
mine, and iodine bear a close resemblance to one another in their 
properties and chemical behavior; collectively they are called 
the halogens (from halite, the scientific name for rock salt). 
In the present chapter we shall first briefly review what has 
already been learned about chlorine and some of its compounds, 
and then after a more extensive consideration of the chemistry 
of chlorine take up a study of the remaining members of this 
important group of elements. 

232. Resume of Facts Already Learned. — We know that 
common salt, NaCl, is the most abundant compound of chlorine; 
it forms the raw material from which all other compounds of 
chlorine as well as the free element are made. The action of 
sulfuric acid on salt (103) yields hydrochloric acid which, 
by electrolysis (43) or by the action of lead dioxide, gives free 
chlorine (167). With bases or metallic oxides hydrochloric acid 
yields chlorides, as illustrated by the following reactions : 

K0H+HC1^KC1+H 2 (107) 

MgO+ 2 HCl->MgCl 2 +H 2 0. (143) 

Chlorides also result when carbonates are treated with hydro- 
chloric acid (163): 

Na 2 C0 3 +2HCl->2NaCl+C0 2 +H 2 
CaC0 3 +2HCl->CaCl 2 +C0 2 +H 2 0. 

It will be recalled that the chlorides of silver, lead, and 
univalent mercury are almost insoluble in water (167, 169, 182); 
these salts are easily obtained by the action of solutions of hydro- 
chloric acid or any soluble chloride on solutions of soluble salts 
of these metals, thus: 

AgN0 3 +HCl->AgCl+HN0 3 

Pb(N0 3 ) 2 +2NaCl->PbCl 2 +2NaN0 3 . 

13s 



136 Introduction to General Chemistry 

The metals which react with hydrochloric acid set free hydrogen 
and are themselves converted into chlorides, for example: 

Zn+2HCl->ZnCl 2 +H 2 . (149) 

2A1+6HC1->2A1C1 3 + 3 H 2 . (174) 

Chlorides also result from the direct union of chlorine with other 
elements : 

H 2 +C1 2 -^ 2 HC1 (44) 

2Al+3Cl a ->2AlCl 3 . (174) 

233. The Occurrence of Chlorine Compounds in Nature. — ■ 
-Free chlorine does not occur in nature. If free chlorine were 
present in nature it would very soon unite with other substances 
to form compounds. Common salt is by far the most abundant 
natural compound of the element. It occurs as a mineral, rock 
salt (halite), and as dissolved salt in sea-water and the waters of 
salt lakes and springs. Sea-water contains about 3 per cent, 
while the water of Great Salt Lake in Utah contains about 
20 per cent, of salt. Rock salt has doubtless been formed in 
past geological times by the slow, natural evaporation of sea- 
water. Other chlorides, particularly those of potassium, KC1; 
magnesium, MgCl 2 ; silver, AgCl; and lead, PbCl 2 , are also found 
in nature. 

234. The Discovery of Chlorine. — Free chlorine was first 
made by the Swedish chemist Scheele, in 1774, and therefore 
practically at the same time that Lavoisier in France discovered 
the true explanation of burning. Scheele made chlorine by 
the action of hydrochloric acid on manganese dioxide, a mineral 
having the formula Mn0 2 , and therefore an oxide of the metallic 
element manganese. The reaction occurs thus: 

4 HCl+Mn0 2 ->MnCl 2 +Cl 2 +2H 2 0. 

Chlorine was not thought to be an element until nearly forty 
years after its discovery, but was believed to be an oxide of 
hydrochloric acid, until a famous English chemist, Sir Humphrey 
Davy, showed by conclusive experiments that it did not contain 
oxygen and was really an elementary substance. 



Halogens with Hydrogen and Metals 



137 



235. The Preparation of Chlorine from Hydrochloric Acid. — 
We have already seen (167) that chlorine is formed when lead 
dioxide is warmed with hydrochloric acid : 

4 HCl+Pb0 2 ->PbCl 2 +Cl 2 +2H 2 0. 

This reaction is entirely analogous to the one between hydro- 
chloric acid and manganese dioxide mentioned in the preceding 
paragraph, and since the last substance is cheaper than lead 
dioxide it is the one commonly used in the laboratory for the 
preparation of chlorine. The experimental method consists in 
adding to, say, 100 g. of granular manganese dioxide contained 
in a flask about 300 c.c. of concentrated 
hydrochloric acid and warming gently: 

4 HCl+Mn0 2 ^MnCl 2 +Cl 2 +2H 2 0. 

Manganese chloride is an easily soluble 
salt which forms pink crystals of a 
hydrate, MnCl 2 • 4.H 2 0. 

An excellent, though expensive, 
method of making small amounts of 
chlorine for experimental work in the 
laboratory consists in allowing concentrated hydrochloric acid to 
drop slowly onto solid potassium permanganate, KMn0 4 (Fig. 30). 
The latter substance is one of the most powerful oxidizing agents 
and reacts rapidly in the cold with hydrochloric acid, thus : 

i6HCl+2KMn0 4 ^2KCl+2MnCl 2 +5Cl 2 +8H 2 0. 

Since the rate of production of chlorine is easily regulated by 
control of the rate of flow of the acid, the method is a very con- 
venient one for the lecture table. 

236. Chlorine, a Poisonous Gas.— The chlorine which is given 
off is a heavy, yellowish, poisonous gas having an exceedingly 
violent action on all mucous membranes. It is the gas which 
was first used with such frightful effect in the trenches in the 
European war. Great care must be exercised to prevent the escape 
of appreciable amounts of chlorine into the air of the laboratory 




Fig. 30 



138 Introduction to General Chemistry 

and to avoid as far as possible inhalation of the gas. Waste 
chlorine is easily absorbed when passed into a solution of 
caustic soda. 

237. The Electrolytic Preparation of Chlorine. — We have 
already learned ^43^ that chlorine is formed when hydrochloric 
acid is electrolyzed. By means of the Brownlee apparatus 
shown in Fig. 21 it is found that equal volumes of hydrogen 
and chlorine are formed when the concentrated acid is used. 
If, however, very dilute acid is used, then the products are 
largely hydrogen and oxygen formed by the decomposition of 
the water, and very little chlorine is set free. A complete 
explanation of this curious fact is not possible until certain 
matters treated in a following chapter have been considered; 
but it may be stated that hydrochloric acid is more easily 
decomposed than water by the electric current, and that if much 
of the former is present in a water solution it is decomposed by 
preference to the water. In the electrolysis apparatus the poles 
or electrodes are sticks of carbon. The hydrogen is liberated 
at the negative pole, the chlorine at the positive pole. 

238. The Electrolysis of Common Salt. — The electrolysis of 
a concentrated solution of common salt is by far the most im- 
portant practical method for the manufacture of chlorine. It 
is a process which is carried out on a very large scale, as at 
Niagara Falls, where electrical power is cheap and yields not only 
chlorine but also hydrogen and caustic soda. We might expect 
the products of the electrolysis of salt to be sodium and chlorine, 

2NaCl-^2Na+Cl aj 

but when we recall that sodium reacts at once with water to 
form hydrogen and sodium hydroxide (caustic soda), the actual 
result appears reasonable. A more complete explanation must 
be deferred until later. As in the case of the electrolysis of 
hydrochloric acid the chlorine is set free at the positive electrode, 
which is a carbon plate, while the sodium and hydrogen are 
formed at the negative electrode. 

239. Deacon's Process. — Before the electrical method just 
described was used practically, a process invented by Deacon 



Halogens with Hydrogen and Metals 139 

was the cheapest technical method of making chlorine. This 
process is based on the fact that a mixture of hydrogen chloride 
gas and oxygen react at a high temperature to form chlorine and 

water. 

4HC1+0 2 ->2C1 2 -2H/j. 

This reaction scarcely takes place at all at ordinary temperatures, 
and even at the most favorable high temperature it takes place 
very slowly. Deacon discovered that the reaction could be 
greatly hastened if the heated mixture of hydrogen chloride and 
oxygen were passed over broken bricks coated with copper 
chloride,. CuCl 2 . A small amount of this substance is able to 
promote the reaction of almost unlimited amounts of the reacting 
gases without itself being permanently changed or desti 
A substance that behaves in this way is called a catalytic agent. 
Catalytic agents of various sorts are extensively employed in 
chemistry. In the Deacon process air. which is essentially a 
mixture of oxygen and nitrogen, may be used instead of pure 
oxygen, which would be too expensive for practical purpc 

240. A Remarkable Phenomenon: Chemical Equilibrium. — 
It is a remarkable fact that even under the most favorable con- 
ditions the reaction between hydrogen chloride and oxygen does 
not go to completion, but stops while the gaseous mixture still 
contains some of both of these gases. The cause is discovered 
when we find that steam and chlorine react at about 400' to give 
some hydrogen chloride and oxygen: 

2CL-2H : 0-^_iHC:-0 2 . 

This is. in fad. exactly the reverse of the reaction we ha'ce been con- 
sidering. It is plain, therefore, that the failure of the reaction 
between hydrogen chloride and oxygen to go to completion is 
due to the interaction of the products, chlorine and water, to 
form again some of the first pair of gases. 

If a mixture of hydrogen chloride and oxygen in the propor- 
tions shown in the equation is heated to a constant temperature, 
say 400', a mixture finally results in which all jour of the sub- 
stances are present in definite proportions. A mixture having 
exactly the same proportions of each of the four substances 



140 



Introduction to General Chemistry 




Fig. 31 



results if the starting substances are chlorine and water, taken 
also in the proportions indicated by the equation. In the 
mixture which finally results, the four substances are said to be 
in a state of chemical equilibrium. The subject of chemical 
equilibrium is a very important one which is to be studied in 
detail in the next chapter. 

241. The Physical Properties of Chlorine. — Chlorine is a 
pale-yellow gas, having a density about two and a half times as 
great as air. Under standard conditions one liter weighs 3 . 22 g. 
Chlorine is rather soluble in water, 100 c.c. of water at 20 dis- 
solving 226 c.c. of the gas. For this reason the gas is not easily 

collected over water; on account 
of its high density it is easily 
collected by the downward dis- 
placement of air. If a water 
solution of chlorine is cooled 
nearly to o°, yellow crystalline 
chlorine hydrate, having the 
formula C1 2 -8H 2 0, is formed. 
This hydrate is very unstable and decomposes slowly at room 
temperature and rapidly at higher temperatures into chlorine 
gas and water. 

242. The Liquefaction of Chlorine. — A very interesting and 
important experiment was once made with this hydrate by the 
great English physicist and chemist Faraday, who was at the 
time assistant to Sir Humphrey Davy (234). Crystals of 
chlorine hydrate were sealed up in one end of a bent glass tube, 
as shown in Fig. 31; when the hydrate was gently warmed 
while the other end of the tube was cooled with ice a yellow 
liquid formed in the cold end of the tube. This liquid proved 
to be liquefied chlorine. It is a heavy, mobile liquid, which is 
easily obtained from chlorine gas either by cooling the latter 
to about 40 below zero at atmospheric pressure, or by com- 
pressing it to about four atmospheres' pressure at about o°. 
Under one-atmosphere pressure liquid chlorine boils at — 34 . 
This work of Faraday in liquefying chlorine was of very great 
importance, since it was the beginning of the epoch-making 



Halogens with Hydrogen and Metals 



141 



experiments in which he succeeded in liquefying all known 
gases except five, among which were hydrogen, oxygen, #nd 
nitrogen. 

243. The Union of Chlorine and Hydrogen. — Chlorine and 
hydrogen do not react at an appreciable rate at room tempera- 
ture if kept in complete darkness, but do unite with explosive 
violence if exposed to a bright light, hydrogen chloride being 

formed, thus: 

H 2 +C1 2 ->2HC1. • 

In order to demonstrate this interesting phenomenon a thin- 
walled glass bulb is filled with a mixture of equal volumes of 
the two gases ; the bulb is then covered with a thick-walled bell 
jar (Fig. 32) and strongly illuminated either by direct sunlight 




m 



fT 



ffi 



Fig. 32 



Fig. 33 



or by the rays from burning magnesium ribbon. The sharp 
explosion which follows reduces the glass bulb to a powder, but 
does no damage to the bell jar. The mixture of chlorine and 
hydrogen is best obtained by the electrolysis of concentrated 
hydrochloric acid in the apparatus shown in Fig. 33. The inner 
vessel has two carbon electrodes. It is surrounded by a larger 
vessel, through which water flows to prevent rise of temperature. 
During the filling of the bulb and up to the time all is ready for 
the explosion it must be shielded from bright light. The union of 
chlorine with hydrogen takes place slowly, without explosion, if 
the mixture of the two gases is exposed for a sufficient length of 
time to moderate light (44). 

244. The Burning of Hydrogen in Chlorine. — If a jet of hydro- 
gen burning in air is lowered into a jar of chlorine it continues 



142 



Introduction to General Chemistry 



r 



to burn with a pale flame (Fig. 34). The flame is the result of 

the intense heat produced by the union of the two gases to form 

hydrogen chloride. 

245. The Action of Chlorine on Water. — Water dissolves 

about two or three times its own volume of chlorine at room 

^ temperature, giving a yellowish solution 

known as chlorine water. This solution 
smells strongly of chlorine and is often 
used in the laboratory in place of 
chlorine gas. If chlorine water is 
exposed to light it soon loses its color 
and odor, and at the same time a color- 
less, odorless gas, which proves to be 
oxygen, is given off. The experiment 
may readily be carried out in the 
manner shown in Fig. 35. A cylinder 
rilled with chlorine water is inverted in 

a dish or beaker and exposed to bright light for a day or two. 

The gas produced will be found to be oxygen, formed according 

to the equation 

2Cl 2 + 2 H 2 0^ 4 HCl+0 2 . 






Fig. 34 



This is the reversal of the reaction by which chlorine is made by 
Deacon's process. While chlorine gas and steam react only 
partially at a high temperature, as already 
stated, chlorine dissolved in water and ex- 
posed to light reacts slowly, hut completely, at 
room temperature to form hydrochloric acid 
and oxygen. This curious difference in be- 
havior may be traced to the fact that while 
gaseous hydrogen chloride and oxygen react 
to the extent of about 80 per cent at 400 , 
oxygen gas does not act at all on a solution of 
hydrochloric acid at -room temperature. No chlorine and water, 
therefore, can be reproduced in cold water solution from the 
products of the action of these two substances, and so the 
main reaction goes on to completion. Much more is known 



Fig. 35 



Halogens with Hydrogen and Metals 143 

about the action of chlorine on water than is contained in this 
paragraph, and the subject will be taken up again in the following 
chapter. 

246. The Union of Chlorine with Metals. — Chlorine unites 
directly with many metals forming chlorides. In many cases 
the reaction takes place at once, with the production of heat and 
even in some cases of light, upon bringing the metal into chlorine 
gas. Thin pieces of copper in the form of dutch metal take fire 
when dropped into a jar of chlorine, forming copper chloride, 

Cu+Cl 2 ->CuCl 2 . 

The metal antimony (symbol Sb), in the form of powder, also 
unites with chlorine, with the production of light and heat, if 
sifted into a cylinder of the gas, antimony trichloride being 

formed, 

2Sb+3Cl 2 ^2SbCl 3 . 

Chlorine also unites directly with sodium, potassium, 
magnesium, zinc, iron, aluminum, mercury, and many other 
metals to form the corresponding chlorides. 

247. The Union of Chlorine and Phosphorus. — The element 
phosphorus is a white, waxy solid which can be made from cal- 
cium phosphate, bone ash (158). We have already seen (10) 
that phosphorus burns readily in the air. In so doing it unites 

with oxygen, thus: 

4 P+50 2 ->2P 2 5 , 

forming a white, solid product, phosphorus pentoxide. Phos- 
phorus also unites directly with chlorine to form either phos- 
phorus trichloride, PC1 3 , or phosphorus pentachloride, PC1 5 . 

The preparation of the trichloride may be carried out in a retort 
as shown in Fig. 36. About 20 g. of dry phosphorus are placed 
in the retort and a stream of chlorine, dried by passing it through 
a wash bottle containing concentrated sulfuric acid, is passed 
in by means of the glass tube which passes through the stopper 
of the retort. As soon as the chlorine reaches the phosphorus, 
union takes place with the formation of much heat and the appear- 
ance of a pale flame. The course of the reaction is readily con- 
trolled by regulating the rate of flow of the gas and by moving 



144 



Introduction to General Chemistry 



the gas inlet tube up or down in the retort. If the contents get 
too hot so that phosphorus begins to distil, the temperature can 
be lowered by raising the tube. On the other hand, if yellowish 
crystals of the pentachloride appear in the retort, the tempera- 
ture is too low and the tube should be lowered. The reaction 
occurs thus: 

2P+ 3 C1 2 ->2PC1 3 . 

Phosphorus trichloride distils over and condenses to a liquid 
in the cooled receiver. It may be purified by being distilled 




Fig. 36 

from a clean, dry retort. It is a colorless liquid which boils at 
74 . It readily unites with more chlorine, forming solid crystal- 
line pentachloride, PC1 S : 

PC1 3 +C1 2 ->PC1 S . 

The chlorides of phosphorus are not salts. Both compounds 
are acted upon vigorously by water, according to the following 

equations: 

PC1 3 + 3 H 2 0->H 3 P0 3 + 3 HC1 
PC1 5 + 4 H 2 0->H 3 P0 4 +5HC1. 



The products are hydrochloric acid and in the first case phos- 
phorous acid, H 3 P0 3 , and in the second case phosphoric acid, 



H 3 PO, 



(i59). 



Halogens with Hydrogen and Metals 145 

248. Chlorine and Turpentine. — Turpentine is a colorless 
liquid having the formula C I0 H l6 . It reacts violently with 
chlorine, thus: 

C I0 H i6 +8C1 2 ->ioC+i6HC1. 

The reaction is best shown by bringing a strip of filter paper 
which has been dipped in turpentine into a cylinder of chlorine; 
a flash of flame occurs accompanied by a dense, black smoke, due 
to the finely divided carbon formed. This reaction, as well as 
that between chlorine and water, shows the great tendency of 
chlorine to unite with hydrogen even if the hydrogen is in the form 
of a compound. 

249. Practical Uses of Chlorine. — A piece of litmus paper 
dipped into chlorine water becomes colorless. Many other 
vegetable colors are also bleached in the same way. The process 
is of great practical importance. All white cotton goods have 
been bleached by a modification of this process, which will be 
described in another chapter (351). 

In recent years a new and important use for chlorine has 
been found as a reagent for the sterilization of municipal water 
supplies. The effectiveness of chlorine is due to the fact that 
it is a powerful germicide by reason of its great chemical activity. 
The chlorine is dissolved in the water at the pumping stations 
and during the interval required for the water to flow through 
the mains it reacts with the germs present and is itself reduced 
to harmless chlorides. The water supply of the city of Chicago 
is purified in this way. 

250. The Preparation of Hydrochloric Acid. — We have 
already learned that hydrogen chloride is made by the action 
of sulfuric acid on common salt. The best laboratory method 
is that described earlier (103), the reaction taking place according 
to the following equation: 

NaCl+H 2 S0 4 ->NaHS0 4 +HCl. 

If, however, double the proportion of salt indicated by this equa- 
tion is taken and the temperature is finally raised sufficiently, 
the following reaction will take place: 

2NaCl+H 2 S0 4 ->Na 2 S0 4 + 2HCI. 



146 Introduction to General Chemistry 

By the last reaction a given quantity of sulfuric acid will produce 
double the quantity of hydrogen chloride as in the first; it is 
therefore the more economical and is the one used in the com- 
mercial production of hydrochloric acid. 

The union of hydrogen and chlorine to form hydrogen 
chloride has already been discussed (44, 243). In recent years, 
since chlorine has become available in immense quantities as 
a by-product of the manufacture of caustic soda, some hydro- 
chloric acid has been produced commercially in this way. 

The old name for hydrochloric acid was muriatic acid, and 
this is the name by which the crude acid is still commonly known 
in trade. 

251. The Physical Properties of Hydrogen Chloride, — 
Hydrogen chloride is a colorless gas, having a choking odor and 
forming a cloud of white fumes in moist air. Its density is 
somewhat greater than that of air; one liter weighs 1.642 g. 
The gas is very soluble in water; at room temperature water 
dissolves about 450 times its volume of the gas, giving a con- 
centrated solution of hydrochloric acid. Considerable heat is 
produced when the gas dissolves in water, so that the solution 
becomes decidedly warm. In general, when gases dissolve in 
water heat is produced. So-called chemically pure hydrochloric 
acid has a specific gravity of 1 . 2 and contains about 37 per cent 
of hydrogen chloride, the balance being water. 

When the 37 per cent acid is heated, hydrogen chloride 
gas is given off and the remaining solution becomes less concen- 
trated. Upon continued heating in an open vessel, the tempera- 
ture rises to uo° before the liquid boils; by this time the 
concentration has decreased to 20 per cent. As the solution 
continues to boil, its concentration, 20 per cent, and boiling- 
point, no°, remain constant; the condensed vapor, the so-called 
distillate, also has a concentration of 20 per cent. 

On the other hand, if very dilute hydrochloric acid is boiled 
it loses water chiefly and becomes more concentrated; finally, 
when the concentration has reached 20 per cent the boiling 
temperature has become no°, after which both concentration 
and boiling-point remain constant. 



Halogens with Hydrogen and Metals 147 

252. The Chemical Properties of Hydrochloric Acid.— The 
most important chemical properties of hydrochloric acid have 
already been studied. These may be briefly reviewed in this 
paragraph. Hydrochloric acid is perhaps the most typical of 
all acids; it turns litmus red and its very dilute solution, say 
1 per cent, has a pleasant sour taste ; it neutralizes the hydroxides 
and oxides of metals, forming chlorides and water, for example: 

NaOH+HCl->NaCl+H 2 
CuO+ 2HCl->CuCl 2 +H 2 0. 

It acts on many metals forming chlorides and hydrogen, thus : 
Fe+2HCl->FeCl 2 +H 2 . 

The addition of hydrochloric acid to solutions of salts of 
silver (169), lead (167), and univalent mercury (182) gives 
precipitates of insoluble chlorides, thus : 

AgN0 3 +HCl->AgCl+HN0 3 . 

Oxidizing agents, such as oxygen gas at a high temperature 
and higher oxides of the metals like manganese dioxide, liberate 

chlorine : 

4 HC1+0 2 ^ 2 C1 2 +2H 2 (239) 

4 HCl+Mn0 2 ->Cl 2 +MnCl 2 H-2H 2 0. (235) 

Hydrochloric acid is an almost indispensable chemical 
reagent. It is used extensively both in scientific and in technical 
work. It is manufactured in large quantities and is an impor- 
tant article of commerce. 

253. The Action of Hydrochloric Acid on Sodium Hydrogen 
Sulfate. — If concentrated hydrochloric acid is added slowly, 
with stirring, to a concentrated solution of sodium hydrogen 
sulfate, a white crystalline precipitate is formed, which, when 
filtered out, washed with a little water, and dried, is found to 
consist of pure sodium chloride. The reaction is represented 
thus: 

HCl+NaHS0 4 ^NaCl+H 2 S0 4 . 

This is seen to be just the reverse of the reaction by which 
hydrogen chloride is made from salt. It is therefore a reversible 



148 Introduction to General Chemistry 

reaction. The direction which the reaction will take depends 
upon the amount of water present and the temperature. Dry 
salt and anhydrous (water-free) sulfuric acid react practically 
completely to form hydrogen chloride and sodium hydrogen 
sulfate; while sufficiently dilute sulfuric acid and salt do not 
give off any hydrogen chloride gas. The reason is simple: the 
gas is very soluble in water, and even if it were formed it would 
remain dissolved in the water present. The fact that con- 
centrated solutions of hydrogen chloride and sodium hydrogen 
sulfate gives a precipitate of solid sodium chloride shows clearly 
that the reaction has a reversible tendency. It seems probable 
that in the presence of much water, that is, in dilute solution, 
all four of the substances are present in any solution that is made 
by bringing either pair of substances together. In such a solution 
we may say that there exists a state of equilibrium as the result 
of each pair of substances on the same side of the equation con- 
tinuously reacting to form the pair on the opposite side, thus: 

H 2 S0 4 +NaCl±>NaHS0 4 +HCL 

254. Bromine. — The element bromine (symbol Br) resembles 
chlorine more closely than does any other element. It does not 
occur free in nature. Its salts, the bromides, are frequently 
found in small amounts associated with chlorides. Sea-water 
contains a small proportion of bromides. Large quantities of 
bromides are obtained from deposits accompanying those of 
sodium nitrate in the desert regions of Chile. The brines from 
salt springs in Michigan also furnish bromides in commercial 
quantities. 

255. Sodium bromide, NaBr, potassium bromide, KBr, and 

magnesium bromide, MgBr 2 , are the commonest salts directly 

obtainable from natural salt deposits and brines. From any 

of these the element is readily set free by the action of chlorine, 

thus: 

2 KBr+Cl 2 ->2KCl-fBr 2 . 

Upon passing chlorine gas into a solution of potassium bromide, 
the solution turns brown and when heated gives off reddish-brown 
vapors of bromine, which when cooled condense to liquid 



Halogens with Hydrogen and Metals 149 

bromine. Bromine is a reddish-brown liquid which has a 
density over three times that of water. It boils at 58 and readily 
volatilizes at ordinary temperatures. The vapor is, if anything, 
more irritating to mucous membranes than chlorine, and the 
liquid produces deep burns when brought into contact with the 
skin. Bromine must be handled with extreme caution. In case 
of accident wash off the bromine with water immediately; then 
consult an instructor regarding further treatment. 

Bromine dissolves in water to the extent of about 3 per cent 
to form a light-brown solution, known as bromine water. 

256. Hydrobromic Acid, HBr. — Hydrogen bromide, the 
water solution of which is known as hydrobromic acid, can be 
made by the direct union of its constituent elements: 

H 2 +Br 2 -^ 2 HBr. 

The best method of making hydrogen bromine is based on the fact 
that bromine unites with phosphorus to form a tribromide or a 

pentabromide, thus, 

P+ 3 Br->PBr 3 
P+5Br->PBr 5 . 

These compounds are entirely analogous to PC1 3 and PC1 5 (247). 
The bromides of phosphorus also resemble the chlorides in their 
reactions with water, thus : 

PBr 3 + 3 H 2 0->H 3 P0 3 +3HBr 
PBr 5 + 4 H 2 O^H 3 P0 4 +5HBr. 

The preparation of hydrobromic acid is carried out in the 
apparatus shown in Fig. 37. 

Ten grams of red phosphorus, 
10 c.c. of water, and 20 to 25 g. 
of quartz sand are placed in a 
250 c.c. flask and 15 c.c. of bro- 
mine, contained in the dropping 
funnel, are allowed to run in 
slowly, drop by drop. The U-tube 
contains some pieces of broken 
glass or brick or similar inert 
material mixed with 3 or 4 g. of red phosphorus, the object of 



rv 



150 Introduction to General Chemistry 

the glass or brick being to distribute the phosphorus so that it 
will present the maximum of surface. The hydrogen bromide 
.given off is freed from accompanying bromine vapor by the 
phosphorus in the U-tube and is absorbed by water contained 
in the cylinder. The delivery tube should not dip into the 
Water in the cylinder, since the gas is so soluble that there would 
be danger of water getting back into the U-tube and flask. 

257. The Properties of Hydrogen Bromide. — Hydrogen 
bromide is a colorless gas with a choking odor; it gives white 
fumes in moist air and dissolves abundantly in water to form a 
solution known as hydrobromic acid. This is a colorless liquid 
which closely resembles hydrochloric acid in its properties. It 
neutralizes bases and unites with metallic oxides to form salts 
called bromides, for example : 

* NaOH+HBr-^NaBr+H 2 
MgO+ 2HBr->MgBr 2 +H 2 
CuO-j-2HBr->CuBr 2 +H 2 
Al(OH) 3 + 3 HBr->AlBr +3H 2 0. 

The bromides of silver, lead, and univalent mercury are almost 
insoluble in water, as are the chlorides of these same metals (252). 
All other bromides are easily soluble. The addition of hydro- 
bromic acid or any soluble bromide to a solution of a salt of 
silver, lead, or univalent mercury gives a white precipitate of 
the insoluble bromide, thus: 

Pb(N0 3 ) 2 +2HBr->PbBr 2 + 2 HN0 3 . 

258. The Oxidation of Hydrobromic Acid. — Hydrogen bro- 
mide and oxygen gases react when heated to form bromine and 

water, 

4 HBr+0 2 ^2Br 2 + 2 H 2 0. 

This reaction is analogous to that between hydrogen chloride 
and oxygen (239), but takes place far more completely, indi- 
cating that hydrogen bromide is more easily oxidized than hy- 
drogen chloride. Other oxidizing agents, such as manganese 
dioxide, readily set free bromine: 

4 HBr+Mn0 2 ^MnBr 2 +Br 2 + 2 H 2 0. 



Halogens with Hydrogen and Metals 151 

In the technical preparation of bromine by means of this 
reaction sodium bromine is treated with dilute sulfuric acid 
and manganese dioxide. In this case all of the available 
bromine is set free. 

2NaBr+2H 2 S0 4 +Mn0 2 ^Na 2 S0 4 +MnS0 4 +Br 2 +2H 2 0. 

259. The Action of Chlorine on Bromides. — A solution of 

any bromide reacts with chlorine to form a chloride and free 

bromine, 

2 KBr+Cl 2 ->2KCl+Br 2 . 

Similarly, hydrobromic acid and chlorine give hydrochloric acid 
and bromine. These reactions are nearly complete, that is, 
they are not reversible to any marked extent, so that we may 
conclude that the metals and hydrogen form by preference com- 
pounds with chlorine rather than with bromine. This fact may 
also be expressed by saying that chlorine has greater affinity than 
bromine for metals and hydrogen. Using this mode of expres- 
sion, we should also say that oxygen has greater affinity than 
bromine for hydrogen, since hydrogen bromide and oxygen give 
water and free bromine. 

260. The Uses of Bromine and Its Compounds. — Potassium 
and sodium bromides are used extensively in medicine as seda- 
tives. Silver bromide is the light-sensitive substance of photo- 
graphic plates. The free element is extensively used in the 
manufacture of important coal-tar dyes. 

261. Iodine. — The element iodine (symbol I), bears almost 
the same relation to bromine that the latter bears to chlorine. 
It does not occur free in nature, but is readily prepared from 
its compounds, the jodides of sodium or potassium, which are 
obtained from two principal natural sources. 

Certain seaweeds contain small amounts of combined iodine 
which has been taken up from sea-water in which a minute 
quantity is present. The ashes left upon burning the dried 
seaweed yield by extraction with water sodium iodide, Nal, 
and potassium iodide, KI. Iodine compounds are also obtained 
as by-products in the purification of the sodium nitrate found 



152 Introduction to General Chemistry 

in Chile (104). Iodine is set free from iodides by the action of 
chlorine, thus: 

2NaI+Cl 2 ^2NaCl+I 2 . 

It is also liberated by the action of manganese dioxide and 
sulfuric acid. 

262. The Physical Properties of Iodine. — Iodine is an almost 
black, crystalline substance, having a density of nearly five. It 
melts at 114 and boils at a somewhat higher temperature, pro- 
ducing a vapor having a magnificent violet color. At a tempera- 
ture slightly below its melting-point iodine has so great a vapor 
pressure that by cautious heating it may be volatilized com- 
pletely without being melted. If the vapor is allowed to strike 
a cold surface crystals of iodine deposit directly without pre- 
liminary formation of liquid iodine. The sublimation (179) of 
iodine in this way is an important step in the purification of this 
element. 

Iodine is very slightly soluble in water, giving a faintly 
brownish solution. It dissolves abundantly in water solutions 
of potassium or sodium iodide. It dissolves easily in alcohol, 
forming a dark-brown solution called by druggists tincture of 
iodine. Iodine also dissolves easily in ether, forming a brown 
solution, and in chloroform and carbon disulfide, forming 
violet-colored solutions. 

263. Iodine and Starch. — If a dilute solution of iodine is 
added to water containing a little starch paste, made by boiling 
starch with 50 to 100 times its weight of water, a deep blue- 
colored solution results. This reaction is a characteristic and 
very delicate test for free iodine. Iodides, like KI, do not give 
this test; but by adding chlorine to a solution of an iodide the 
element is set free and can then be recognized by the starch test. 
An excess of chlorine interferes with this test. 

264. Hydrogen Iodide, HI. — Iodine and hydrogen unite 
slowly at a temperature of 400 to form hydrogen iodide, thus: 

I 2 +H 2 ^ 2 HI. 

The product is a colorless gas, analogous to hydrogen chloride 
and hydrogen bromide. Like these latter gases it dissolves 
abundantly in water, and forms fumes in moist air. 



Halogens with Hydrogen and Metals 153 

Hydrogen iodide is easily made by a reaction resembling 
that used for making hydrogen bromide. Iodine forms with 
phosphorus a tri-iodide, PI 3 . This reacts with water to form 
phosphorous and hydriodic acids thus : 

Pl3+3H 2 0-»H 3 P0 3 +3HI. 

The process of making hydrogen iodide is carried out by placing 
a mixture of powdered iodine and red phosphorus in a flask and 
running in water, drop by drop from a dropping funnel, care 
being taken not to use more water than is necessary, since an 
excess of water would dissolve the gas and so prevent its escape 
from the flask. The apparatus used for making hydrogen 
bromide, Fig. 37, may be used in this case. The U-tube con- 
taining red phosphorus serves here to remove iodine vapor. 
The hydrogen iodide gas may be collected by downward dis- 
placement of air or it may be dissolved in water to form a solu- 
tion of hydriodic acid. 

265. Hydriodic acid is colorless when pure, but is brown if 
it contains free iodine, which it dissolves readily. It neutralizes 
bases and so yields salts called iodides, for example: 

HI+NaOH^NaI+H 2 
2HI+Ca(OH) 2 ->CaI 2 +2H 2 0. 

Hydriodic acid acts on metals similarly to hydrochloric acid, 
giving iodides and hydrogen, thus: 

2HI+Zn^ZnI 2 +H 2 . 

Hydriodic acid is much more easily oxidized than is hydro- 
bromic acid, which in turn is more easily oxidized than hydro- 
chloric acid; while all three acids are oxidized by powerful 
oxidizing agents such as manganese dioxide and lead dioxide; 
hydriodic acid, even in dilute solution, is oxidized slowly by 
atmospheric oxygen, which has no action whatever on dilute 

hydrochloric acid: 

4 HI+0 2 ->2H 2 0+2l 2 . 

The iodine which is slowly liberated according to the equation 
given above remains dissolved in the unchanged acid and gives 
it a brown color. 



154 Introduction to General Chemistry 

266. Uses of Iodine and Iodides. — Iodine is used extensively 
in certain processes of analysis and also in the preparation of 
important compounds containing the element carbon, so-called 
organic compounds. Iodine in the form of tincture of iodine, 
which is a solution of iodine in alcohol, is used externally as an 
antiseptic and also as a counterirritant in medicine. The iodides 
of potassium, sodium, and ammonium are of great importance 
for internal administration in medicine. 

267. Fluorine. — The element fluorine (symbol F), is classed 
among the halogens, although it is less closely related to the 
other three halogens, chlorine, bromine, and iodine, than these 
three are to one* another. The atomic weights of these elements 
are: fluorine, 19; chlorine, 35.5; bromine, 80; iodine, 127. 
Fluorine has, therefore, the smallest atomic weight of any of 
the halogens. We might expect it to resemble chlorine more 
closely than it does bromine and iodine and, in fact, such is the 
case. It is a pale-yellow gas which is very active chemically and 
never occurs free in nature. Its most abundant natural com- 
pound is calcium fluoride or fluor-spar, CaF 2 . It also occurs as 
cryolite, sodium aluminum fluoride, 3NaF-AlF 3 . These sub- 
stances are salts of hydrofluoric acid. We might expect that 
free fluorine could be made by oxidizing hydrofluoric acid with 
manganese dioxide, thus: 

4 HF+Mn0 2 ->MnF 2 +F 2 +2H 2 0; 

but we find, in fact, that hydrofluoric acid is entirely unacted 
upon by the most powerful oxidizing agents. The free element 
was first made by Moissan, by the electrolysis of anhydrous 
liquid hydrogen fluoride, in which some potassium fluoride, KF, 
was dissolved to make it conduct electricity readily. The pro- 
ducts of the electrolysis were fluorine and hydrogen: 

2 HF->F 2 +H 2 . 

Fluorine is one of the most active of all elements. It rapidly 
attacks glass and also most metals, and it reacts at once with 
water forming hydrofluoric acid and oxygen: 

2F 2 + 2 H 2 0-> 4 HF-f0 2 . 



Halogens with Hydrogen and Metals 155 

The preparation of fluorine is a matter of great difficulty, for 
which reason it is very seldom made. 

268. Hydrogen Fluoride, HF. — Hydrogen fluoride, a gas 
whose water solution is called hydrofluoric acid, is the most 
important compound of fluorine. It is formed by the action of 
concentrated sulfuric acid on powdered calcium fluoride : 

H 2 S0 4 + CaF 2 ->CaS0 4 + 2HF. 

It is a colorless gas with a choking odor. At temperatures of 
ioo° and higher its density shows that the gas has the formula 
HF; at room temperature the density is more than double 
that expected for a gas with the formula HF. This fact leads 
to the conclusion that the single molecules have become asso- 
ciated, probably to form double or triple molecules such as 
H 2 F 2 and H 3 F 3 . Hydrogen fluoride gas is condensed to a liquid 
merely by cooling it with ice; colorless liquid hydrogen fluoride, 
so obtained, boils at 19 . 

269. Hydrofluoric Acid and Its Salts. — A 30 per cent solu- 
tion of hydrofluoric acid is an important article of commerce. 
The acid has several practical uses. These include the etching 
and polishing of glass, the removal of sand from castings, and 
the preparation of its salts and also of hydrofluo silicic acid, 
H 2 SiF 6 . 

Hydrofluoric acid forms with bases salts called fluorides. 
The soluble fluorides are very effective preservatives, since they 
inhibit the growth of bacteria, molds, etc. But their use in 
foodstuffs is prohibited because of their interference with diges- 
tion. 

Ammonium fluoride, NH 4 F, is used as a disinfectant for 
utensils used in breweries. Sodium fluoride, NaF, is extensively 
used as a vermin exterminator for poultry. 

270. The Action of Hydrogen Fluoride on Quartz. — We must 
now digress a little from the subject in hand in order to be able 
fully to understand one of the most interesting reactions of 
hydrogen fluoride. The substance called quartz is the oxide 
of an element silicon (symbol Si) and has the formula Si0 2 . 
Common sand is more or less pure quartz. Glass, which is 



156 Introduction to General Chemistry 

made by melting together sand, sodium carbonate, and slaked 
lime, may be considered a mixture of sodium silicate, Na 2 Si0 3 , 
and calcium silicate, CaSi0 3 . Hydrofluoric acid and quartz 
react very readily to form gaseous silicon fluoride, SiF 4 , and 
water, thus: 

4 HF+Si0 2 ->SiF 4 +2H 2 0. 

This is a very characteristic reaction; none of the other halogen 
acids have any action on quartz. 

Glass, which is almost unaffected by the other halogen acids, 
is rapidly attacked by either hydrogen fluoride gas or hydro- 
fluoric acid solution. The fluorine unites, not only with the 
silicon, as in the case of quartz, forming silicon fluoride, but also 
with the sodium and calcium forming sodium fluoride, NaF, and 
calcium fluoride, CaF 2 , the reactions being: 

Na 2 Si0 3 +6HF->SiF 4 +2NaF+ 3 H 2 
CaSi0 3 +6HF->SiF 4 +CaF 2 + 3 H 2 0. 

The result is that glass dissolves very easily in hydrofluoric 
acid, in consequence of which this acid cannot be kept in glass 
bottles. Parafline and other waxes, which are not attacked, are 
used for bottles for this acid, while larger containing vessels are 
made of lead. 

271. Etching Glass with Hydrogen Fluoride. — The etching of 
glass may be illustrated by coating a glass plate with a thin 
layer of parafline, and after making a design or inscription by 
means of a hard pencil which will cut through the parafline and 
thus expose the surface of the glass, exposing the plate to the 
action of hydrogen fluoride gas. The gas is easily made by 
mixing a few grams of powdered fluor spar with concentrated 
sulfuric acid in a shallow lead dish. The latter is covered with 
the glass plate and set aside for ten or fifteen minutes. Upon 
removing the parafline, the design will be found to have been 
etched upon the glass. 

272. Hydrofluosilicic Acid, H 2 SiF 6 . — Hydrogen fluoride and 
silicon tetranuoride unite readily in the presence of water to 
form a solution of hydrofluosilicic acid: 

2 HF+SiF 4 ->H 2 SiF 6 . 



Halogens with Hydrogen and Metals 157 

The solution is a colorless, odorless liquid which does not attack 
glass appreciably. It has well-characterized acid properties: 
it reddens litmus, has a sour taste, and neutralizes bases to form 
salts. This acid is important technically. It is made, in practice, 
by the action of hydrofluoric acid solution on quartz sand: 

6HF+ Si0 2 ->H 2 SiF 6 + 2H 2 0. 

The acid is used for the preparation of its sodium, magnesium, 
and lead salts. Sodium fluosilicate, Na 2 SiF 6 , is extensively used 
in making white enameled ware and also white, or so-called milk, 
glass. It is remarkable in being one of the very few nearly 
insoluble salts of sodium. It is obtained as a white precipitate 
when solutions of common salt and hydrofluosilicic acid are mixed. 

2NaCl-fH 2 SiF 6 ->Na 2 SiF 6 + 2HCI. 

Magnesium fluosilicate, MgSiF 6 , easily soluble in water, is 
used to harden concrete. Lead fluosilicate, PbSiF 6 , also easily 
soluble in water, is made as an intermediate product in refining 
lead (Betts's process). 



CHAPTER XIII 

CHEMICAL EQUILIBRIUM 

273. Incomplete Physical Processes. — While many physical 
processes are seemingly complete, there are others which stop 
far short of completion. Thus, for example, if a small bulb of 
water is broken in a large closed bottle, evaporation of the 
water will start at once, but will apparently cease as soon as the 
pressure of the vapor reaches a value which is definite for a 
definite temperature, although much liquid may still remain 

("2). 

If we add to some water an equal weight of common salt, 
the latter will at once start to dissolve and will continue to do so 
until the solution has, for a given temperature, a certain definite 
concentration; then, although much solid salt is still present, 
no further increase in concentration will take place (122). 

When water in a closed vessel, which it fills but partially, 
reaches its maximum vapor pressure for a given temperature, 
we believe (201) that for every molecule that passes from liquid 
to vapor there is one that passes from vapor to liquid. We say 
that there is equilibrium between liquid and vapor. We believe 
that a similar condition exists when a solid apparently stops 
dissolving in a solution (207). The apparent state of rest or 
inaction in both cases is very probably one in which two opposing 
actions exactly counteract the effects of each other. 

274. Incomplete Chemical Reactions. — Just as in the case 
of physical processes, there are also some chemical reactions 
that do not go to completion. We have already studied some 
reactions of this kind and must now consider the matter more 
fully, as it is one of great importance. 

The reaction between hydrogen chloride and oxygen at 400 
has been considered (239, 240) under the heading " Deacon's 
Process." It has been pointed out (245) that only 80 per cent 
of the hydrogen chloride is oxidized when a mixture of this gas 

158 



Chemical Equilibrium 159 

is heated with oxygen in the proportion indicated in the following 
equation : 

4 HC1+0 2 ^2C1 2 +2H 2 0. 

On the other hand, when a mixture of two formula weights 
each of chlorine and water is also heated to 400 , 80 per cent of 
the chlorine remains unchanged, while 20 per cent is converted 
into hydrogen chloride. It thus happens that whether we start 
with the pair of gases on the left side of the foregoing equation 
or the pair on the right, taking in each case the amounts indi- 
cated in this equation, there results a mixture of the four gases 
which has exactly the same amount of each gas present in the 
two cases. It is easy to see that the cause of each reaction being 
incomplete is found in the fact that the products of either reac- 
tion again react in the opposite direction. In the mixture of the 
four gases which finally results we say that a state of equilibrium 
exists and that the apparent stopping of further change is really 
the result of the formation of hydrogen chloride and oxygen at 
just the same rate as that at which these two gases change into 
chlorine and water. 

275. Velocity of Chemical Change. — The idea that a state 
of chemical equilibrium is the result of two opposing changes 
which take place continuously at such rates or with such velocities 
that for every molecule of a given substance formed one also 
disappears would imply that chemical changes take place gradually 
and possibly at definite speeds or velocities. 

It is well known that certain reactions, as for example the 
burning of a candle or the action of an acid on a metal, certainly 
do take place gradually. It is not so plain that if the reaction 
takes place between two perfectly mixed gases or between two 
substances completely dissolved as a uniformly mixed solution 
•that time is required for the reaction to take place. Nevertheless 
it is probable that no reaction, even an explosion, however rapid 
it may be, is absolutely instantaneous. 

The speed or velocity of reaction in a uniformly mixed solu- 
tion may be beautifully and convincingly demonstrated by means 
of the following experiment: 



160 Introduction to General Chemistry 

To 800 c.c. of water contained in a flask there is added 25 c.c. 
of starch solution (made by boiling 2 g. of starch with 100 c.c. 
of water) and 15 c.c. of a 3 per cent solution of iodic acid, HI0 3 . 
The solution is then well mixed and 5 c.c. of a 3 per cent solution 
of sulfur dioxide, S0 2 , is added and the contents of the flask 
are at once thoroughly mixed by being shaken. The time of 
adding the sulfur dioxide solution is accurately noted — best with 
a stop watch. No change will be seen in the colorless solution 
for about 60 seconds, then the whole solution will suddenly turn 
deep blue. The result is startling! 

If the experiment is repeated, using the same amounts of 
water and of each of the three solutions, and if the temperature 
is also the same, it will be found that the time required for the 
change to occur is always the same. If, however, we increase 
the amount of sulfur dioxide solution added from 5 c.c. to 
10 c.c, everything else remaining the same, the time required 
for the change will be decreased to about 30 seconds. The 
increased velocity is the result of the increase in concentration of 
the sulfur dioxide. 

276. The Effect of Temperature on Reaction Velocity. — The 
effect on the velocity of increasing the temperature is easily 
shown by starting with water at 25 instead of at 20 , when it 
will be clear that at the higher temperature the velocity is decidedly 
greater. 

277. The Action of Sulfur Dioxide on Iodic Acid. — The 
chemical changes involved in the reaction just described need 
not greatly concern the student at this time, as they are of less 
importance than the main facts of reaction velocity that they 
serve here to illustrate. But as it is only natural to wonder 
what has happened in such a striking experiment, the equations 
for the reactions may now be given. In the first place, sulfur 
dioxide, S0 2 , and water form sulfurous acid, H 2 S0 3 , 

S0 2 +H 2 0->H 2 S0 3 . 

The latter reacts with the iodic acid, forming hydriodic and 
sulfuric acids, thus: 

HI0 3 +3H 2 S0 3 ->HI+3H 2 S0 4 . 



Chemical Equilibrium 161 

But hydriodic acid can also react with iodic acid to form free 
iodine and water, 

HI0 3 +5HI->3l 2 + 3 HA 

and then the iodine set free acts on the starch to produce the 
blue color. Now this reaction between iodic and hydriodic 
acid does not take place until all the sulfurous acid has dis- 
appeared. The time observed for the appearance of the "blue 
color is therefore essentially that required for the complete 
oxidation of the sulfurous acid. 

278. The Kinetic Hypothesis Applied to Reaction Velocity. — 
The application of the kinetic-molecular hypothesis (chap, x) 
leads to a simple and reasonable explanation of reaction velocity. 

Let us suppose that two gases, A and B, can unite to form 

a compound AB, and let the reaction be represented by the 

equation 

A+B->AB. 

Let us also suppose that the reaction takes place rather 
slowly after the two gases have been thoroughly mixed. We 
may now consider what determines the rate at which A and B 
unite. It is obvious that union can occur only when a molecule 
of A comes in contact with a molecule of B. Such collisions 
will frequently occur by- reason of the rapid motion of both 
kinds of molecules. Now as these collisions are matters of 
chance it is very easy to see that if more molecules of one or 
both kinds are brought into a given space the number of collisions 
of A molecules with B molecules will be increased. On the 
other hand, decreasing the number of one or both kinds of 
molecules will surely decrease the possible collisions of A with B 
molecules. 

Probably not every collision of an A with a B molecule will 
result in a union of the two to form AB; but if, on the average, 
a certain definite fraction of the collisions result in union, then 
we can say that the greater the number of A and B molecules 
present in a given volume, say 1 c.c, of the gas mixture, the 
greater will be the number of AB molecules formed per second. 
If we start with a mixture of equal numbers of A and B 



162 Introduction to General Chemistry 

molecules there will be for a definite pressure and temperature a 
certain number of AB molecules formed per second. After 
a short time the number of A and B molecules will have de- 
creased appreciably, so that now fewer AB molecules will be 
formed per second, and as time goes on, owing to continual de- 
crease in the numbers of A and B molecules present, there will 
be fewer and fewer AB molecules formed per second. The 
result will be that the rate of formation of A B molecules will 
be greatest at the start and will gradually decrease, until finally, 
if the reaction is not reversible, all A and B molecules will have 
united. 

279. The Kinetic Hypothesis Applied to Chemical Equilib- 
rium. — Let us next consider, in the light of the kinetic-molecular 
hypothesis, the state of affairs if a reaction between gases is 
reversible. The case of hydrogen chloride and oxygen will serve 
as a good illustration. The equation is 

4 HC1+0 2 ^2C1 2 +2H 2 0. 

This reaction takes place with moderate velocity at 400 , finally 
reaching a state of equilibrium in which all four of the substances 
are present. 

Suppose we bring into a closed vessel at 400 a mixture of 
hydrogen chloride and oxygen in the proportion indicated by the 
equation; that is, four molecules of the first gas to one of the 
second. The reaction will begin at a certain velocity, mole- 
cules of hydrogen chloride and oxygen disappearing by uniting 
to form molecules 'of chlorine and water vapor. As time goes 
on there will be fewer and fewer hydrogen chloride and oxygen 
molecules present, so that the number of each uniting per second 
and also the number of chlorine and water molecules formed per 
second will continuously decrease. On the other hand, the mole- 
cules of chlorine and water which have been formed begin to 
reunite to form hydrogen chloride and oxygen. As the total 
numbers of chlorine and water molecules present will increase 
as time goes on, so the numbers of these molecules which react 
and so disappear per second will also increase. The final result 
will be that in each second there will be just as many molecules 



Chemical Equilibrium 163 

of chlorine and water disappearing as the numbers of each formed. 
The same sort of thing will be true for the hydrogen chloride and 
oxygen — as many molecules of each will finally be produced per 
second as the numbers that disappear. When this condition 
is reached no further change in the number of any of the four 
sorts of molecules wiU take place, although chemical change will 
go on continuously. The system is then in a state of equilibrium. 

We may now take up a study of a number of reversible reac- 
tions which reach a state of equilibrium. 

280. Ferric Chloride and Ammonium Sulfocyanate. — If we 
add to a very dilute solution of ferric chloride, FeCl 3 , which is 
faintly yellow in color, a dilute solution of ammonium sulfo- 
cyanate, NH 4 NCS, which is colorless, a blood-red solution results. 
This red substance is ferric sulfocyanate, Fe(NCS) 3 , which is 
formed thus : 

FeCl 3 +3NH 4 NCS->Fe(NCS) 3 +3NH 4 Cl. 

Let us now consider how we may discover whether this 
reaction is complete when the two substances on the. left-hand 
side of the equation are mixed in the indicated proportion or 
whether a state of equilibrium results. The experiment may be 
carried out on the lecture table in the following manner: 

To 2 liters of water we add 20 c.c. of a decinormal solution of 
ferric chloride and 20 c.c. of a decinormal solution of ammonium 
sulfocyanate, which is just the amount indicated by the equation 
as required for the amount of ferric chloride present. Let us now 
divide the red solution into four equal portions, which we may 
place in four similar cylinders or beakers. Suppose we now add 
to the solution in one of the cylinders 20 c.c. more of ammonium 
sulfocyanate solution. The solution will be seen to become 
deeper red in color, which means that more red ferric sulfo- 
cyanate has been formed. Now this fact may be explained in 
either of two ways: first, that we had by mistake used, in the 
first place, less than the correct proportion of ammonium sulfo- 
cyanate indicated by the equation; or, secondly, that a state of 
equilibrium existed in the solution and that the increased con- 
centration of ammonium sulfocyanate had shifted the equilibrium 
so as to form more ferric sulfocyanate. 



164 Introduction to General Chemistry 

We can test the truth or falsity of the first supposition very 
easily. If the original mixture contained less than the correct 
proportion of ammonium sulfocyanate, then there would be 
an excess of ferric chloride, and the addition of more of this salt 
would not increase the amount of ferric sulfocyanate and so 
increase the depth of red color. Let us add, therefore, 20 c.c. 
more ferric chloride to the solution in the second cylinder. It 
becomes deeper red! This seems to show that we are dealing 
with a condition of equilibrium as indicated by the double arrows 
of the following equation: 

FeCl 3 +3NH 4 NCS^Fe(NCS) 3 +3NH 4 Cl. 

If such is the case, then the addition of ammonium chloride to 
the solution in the third cylinder should cause a partial fading 
of the red color by reason of the partial disappearance of the red 
ferric sulfocyanate. Now this is actually what happens when the 
experiment is carried out, as can be seen by comparison with the 
color of the solution in the fourth cylinder. 

It is clear, therefore, that we have here a case of chemical 
equilibrium in which all four of the substances represented in the 
equation can coexist in the same solution. When we added more 
ammonium sulfocyanate to the solution in the first cylinder we 
increased the number of molecules of this salt and so increased 
the chances of collision of ferric chloride molecules with am- 
monium sulfocyanate molecules and this increased the number 
of ferric sulfocyanate molecules formed per second. This 
caused an increase in the total amount of the latter salt, and 
thus gave rise very quickly to a new state of equilibrium in 
which the proportion of ferric salt in the form of red sulfocyanate 
was greater than at first. 

The addition of more ferric chloride to the solution in the 
second cylinder caused a similar shift of equilibrium for anala- 
gous reasons. It is a general rule that increasing the concentration 
of either of the reacting substances on the same side of an equation 
causes a shift in equilibrium so as to form more of the substances 
on the other side of the equation. This rule is also illustrated by 
the fact that when more ammonium chloride was added to the 



Chemical Equilibrium 165 

solution in the third cylinder the color partially faded; this 
showed that some of the red ferric sulfocyanate had disappeared, 
and thus indicated that more ferric chloride and ammonium 
sulfocyanate had been formed. 

281. Hydrogen and Iodine. — We have already seen (264) 
that hydrogen unites with iodine vapor with appreciable speed 
at about 400 . The reaction is not complete, but reaches a 
state of equilibrium while there are still considerable uncom- 
bined substances present. The equation is 

H 2 +I 2 ±5 2HI. 

That the reaction is reversible is easily shown by heating hydro- 
gen iodide gas, when the purple vapors of iodine appear. If the 
temperature is 370 , equilibrium is reached when one-fifth of the 
hydrogen iodide has dissociated into free iodine and free hydro- 
gen. This means that out of every 1,000 molecules of hydrogen 
iodide taken, 200 have dissociated and 800 remain when the 
state of equilibrium is reached. The equation shows that one 
molecule of hydrogen and one of iodine are formed by the disso- 
ciation of two molecules of hydrogen iodide. Therefore for 
every 200 molecules of the compound dissociated there would 
be formed 100 molecules of hydrogen and 100 of iodine. The 
equilibrium mixture resulting from every 1,000 molecules of 
hydrogen iodide taken consists, therefore, of 800 molecules of 
hydrogen iodide, 100 molecules of hydrogen, and 100 molecules 
of iodine. 

If we bring together in a closed vessel equal numbers of 
molecules of hydrogen and iodine and heat at 370 until equili- 
brium is reached we shall find that for every 500 molecules of 
hydrogen and 500 molecules of iodine taken there result 800 
molecules of hydrogen iodide, 100 molecules of hydrogen, and 
100 of iodine. In other words, just the same proportion as 
would be obtained by starting with pure hydrogen iodide gas. 

282. The Criterion of Equilibrium. — In all cases of reactions 
reaching a condition of equilibrium the resulting mixture has 
the same proportions of all substances, whether we start with 
the substances on one side of the equation or with equivalent 



1 66 Introduction to General Chemistry 

amounts of those on the other side. Therefore, if we wish to 

know whether a given reaction has reached equilibrium we bring 

together the substances which would be the products of the first 

reaction. If the resulting reaction then gives a mixture of the 

same composition as that obtained in the first case we conclude 

that both reactions have reached equilibrium. 

283. Equilibrium Constant. — In the hydrogen and iodine 

reaction 

H 2 +I 2 ^2HI 

the rate of union of hydrogen and iodine, which we may 
call the speed of the reaction from left to right, will depend on 
the numbers of molecules of these two elements present in each 
c.c. It would seem probable that for a fixed number of hydrogen 
molecules per c.c. the speed of union would vary directly as the 
number of iodine molecules, and vice versa; so that this speed 
should be proportional to the product of the number of hydrogen 
molecules N x and the number of iodine molecules N 2 present in 
each c.c. of the gas mixture. That is, the speed of union, S I? is 
proportional to N x times iV 2 ; or, algebraically, 

S I = k 1 XN I XN 2 , 

where k x is a constant proportionality factor. 

On the other hand, the reverse change involves the formation 
of hydrogen and iodine from hydrogen iodide, and we see by 
referring to the equation that two molecules of hydrogen iodide 
must react in order that one molecule of hydrogen and one 
molecule of iodine may be formed. This fact would make it 
seem necessary for two molecules of hydrogen iodide to collide 
in order that the change could occur. If so, increasing the num- 
ber of HI 1 molecules in each c.c. would increase for each mole- 
cule the chances per second of collision and, in fact, doubling 
the number of molecules of this gas per c.c. would increase the 
total number of the chances per second fourfold, etc. In other 
words, the number of collisions per second of HI molecules with 

x It has become customary in chemical literature to use formulae of simple 
substances as abbreviations for the names of these substances; especially in cases 
of frequent repetition. 



Chemical Equilibrium 167 

one another will be proportional to the square of the number of 
molecules of this sort in each c.c. The details of the method of 
arriving at this conclusion need not be considered at present. 
If we call the speed of change of hydrogen iodide into hydrogen 
and iodine S 2 and call the number of HI molecules in 1 c.c. N 3t 
then it is plain that this speed is proportional to N 3 2 , or 

kJ 2 == K 2 1\ 3 } 

where k 2 is a constant proportionality factor. 

Let us now think of the state of affairs when equilibrium has 
resulted. The speed of formation of hydrogen iodide which 
is equal to the speed of union of hydrogen and iodine, S Iy is now 
just equal to the speed of dissociation, S 2 , of the hydrogen iodide. 
This must be the case, as otherwise further changes in the pro- 
portions of the three substances would still be taking place and 
the mixture would not be in equilibrium.- For the state of 
equilibrium, therefore, we may write 

and hence 

k l XN l XN 2 =k 2 XN 5 2 

or < 

N 3 2 _h 
N Z XN 2 k 2 ' 

Now &! and k 2 are both constant quantities for the reaction 
under consideration if the temperature is fixed, and therefore 
their quotient is a constant, so that we may write 



a constant. Therefore 






N * =K. 



N,XN 



This algebraic equation means that for the condition of equilibrium 
at a fixed temperature the square of the number of molecules per c.c. 
of EI divided by the product of the numbers of molecules of H 2 and 
1 2 is a fixed or constant quantity. This matter can perhaps be 
made a little plainer by use of a numerical example. We have 



1 68 Introduction to General Chemistry 

seen that at 370 the equilibrium mixture which results from 
1,000 original HI molecules consists of 800 molecules of HI, 100 
of H 2 , and 100 of I 2 . In each c.c. of such an equilibrium mixture 
the total number of molecules will be very great; but, of course, 
the numbers of each kind will be in the same proportion as for a 
total of 1,000 molecules, and therefore 

100X100 

If we start with unequal instead of equal numbers of molecules 
of hydrogen and iodine we can calculate by means of the equa- 
tion 

64 



N 3 > 



N X XN 2 

what the state of equilibrium will be. For example, suppose we 
start with a mixture of hydrogen and iodine containing four 
times as much hydrogen as would theoretically be necessary 
for the iodine taken; that is, four molecules of hydrogen for one 
of iodine. Calculation shows that if we start with 800 mole- 
cules of hydrogen and 200 molecules of iodine, when equilibrium 
is reached, out of a total of 1,000 molecules 392 will be hydrogen 
iodide, 604 will be free hydrogen, and 4 will be free iodine. 

284. Ammonia and Water. — Several reactions already studied 
reach a condition of equilibrium; three of the most familiar of 
these may now be considered as additional examples of the subject 
under discussion. Ammonia gas, NH 3 , dissolves abundantly in 
water, giving a solution which turns litmus blue and forms salts 
with acids. The solution contains ammonium hydroxide, 
formed by the union of ammonia with water (91) : 

NH 3 +H 2 0->NH 4 OH. 

The solution smells strongly of ammonia and. if it is boiled a 
short time all of the gas is given off. This shows that ammonium 
hydroxide easily dissociates into its constituents. It seems 
highly probable that in the water solution a condition of equilib- 
rium exists, as indicated in the equation 

NH 3 +H 2 0±?NH 4 OH, 



Chemical Equilibrium 



169 



both free ammonia and ammonium hydroxide being present. 
Heating such a solution renders the free ammonia less soluble, 
and as this partially escapes, the rate of formation of ammonium 
hydroxide falls farther and farther behind the rate of dissociation 
of this compound until finally all of the latter has disappeared. 

285. Carbon Dioxide and Water. — A water solution of carbon 

dioxide, C0 2 , contains carbonic acid, H 2 C0 3 . But such a solution 

easily gives off carbon dioxide, especially if warmed; which leads 

us to conclude that the reaction is a reversible one, and that in 

the solution there is a state of equilibrium as represented by 

the equation 

C0 2 +H 2 0±5H 2 C0 3 . (152) 

286. Sulfur Dioxide and Water. — Sulfur dioxide, S0 2 , which 
is formed when sulfur burns, is a colorless gas with a suffocating 
odor: 

s+o 2 ->so 2 . 

It is easily soluble in water, giving a solution which smells 
strongly of the gas and has acid properties. The solution con- 
tains a compound, sulfur ous acid, H 2 S0 3 . This solution gives off 
all of its sulfur dioxide when boiled, and we conclude, therefore, 
that the acid easily decomposes into its constituents, water and 
sulfur dioxide, and that in the solution we have a state of equilib- 
brium, as represented by the equation 

S0 2 +H 2 O^H 2 S0 3 . 

287^ The Effect of Pressure on a System in Equilibrium. — 

Suppose we have, say, 1 liter of water saturated 
with a gas, say oxygen, at a fixed temperature 
and at one-atmosphere pressure. To say that 
the water is saturated with the gas means that 
a condition of equilibrium exists between solu- 
tion and gas. Let us suppose the solution and 
gas are contained in a cylinder fitted with a gas- 
tight piston (Fig. 38) and that the volume of the 
undissolved oxygen gas above the solution is 1 
liter. If now we double the pressure on the gas 
more of the gas passes into solution, finally producing a new 



JZZ ■\\'////////////A 



Fig. 38 



170 Introduction to General Chemistry 

state of equilibrium. By reason of the fact that part of the 
gas dissolved when the pressure was doubled the volume of the 
remaining gas will not be half a liter, as we should expect if no 
additional quantity of oxygen dissolved in the water present, 
but appreciably less than half a liter. The effect, therefore, of 
increasing the pressure on the system is to cause its volume to 
become smaller than would be the case if no shift of equilibrium 
had occurred. This is the way in which an increase of pressure 
always affects a system in equilibrium : the state of equilibrium 
shifts in such a way as to cause a greater decrease in volume than 
would be the case if no change in the state of equilibrium occurred. 
Let us consider another case. We may inquire how the 
equilibrium represented by the equation 

4 HC1+0 2 ^2C1 2 +2H 2 

would be affected by increase of pressure. We see by reference 
to the equation that four volumes of HC1 and one of 2 give two 
volumes of Cl 2 and two of H 2 0; that is, that when the reaction 
takes place from left to right there is a decrease in volume from 
5 to 4. We should expect, therefore, that by increasing the 
pressure the equilibrium would shift somewhat from left to 
right; that is, that more chlorine and water would be formed 
at the expense of the hydrogen chloride and oxygen; and this 
is exactly what actually happens. 

The effect of increase of pressure on any system in equilibrium 
is, in all cases, to shift the equilibrium so as to favor the formation 
of substances occupying a smaller volume. In case no change of 
volume accompanies a chemical reaction, then the state of 
equilibrium is not affected by change of pressure. The reaction 

H 2 -f-I 2 ±5 2HI 

is an example of this sort. Here one volume of hydrogen and 
one volume of iodine vapor react to form two volumes of hydro- 
gen iodide, so that no change of volume occurs when the reaction 
takes place. It has been found by careful investigation that 
the equilibrium proportion of the three substances is not changed 



Chemical Equilibrium . 171 

by altering the pressure, as long as the temperature remains 
constant. 

288. Effect of Temperature on a System in Equilibrium. — We 
have already learned (112) that the vapor pressure of water 
increases with increase of temperature. We know also that a 
large amount of heat is absorbed when water is evaporated; at 
ioo° it requires 540 calories to change one gram of water into 
steam. This is the so-called latent heat, of vaporization. If 
we have, in a closed vessel, water in equilibrium with its vapor, 
and then apply heat, two effects are produced : the temperature 
is raised and the vapor pressure is increased. The increase in 
pressure is caused by the evaporation of some water, and this 
evaporation absorbs some of the heat which has been applied. 
This is a typical case, for we always rind that when we apply heat 
to any system in equilibrium that the state of equilibrium shifts 
in such a way that heat is absorbed in the change. As heat is 
absorbed when water evaporates, heating causes increased vapor 
pressure. 

The effect of temperature on the solubility of substances 
has already been studied (134). We have learned that heat is 
either absorbed or produced when a substance dissolves; this is 
the so-called heat of solution. Substances which dissolve with 
absorption of heat become more soluble with rise of tempera- 
ture, while those which dissolve with evolution of heat, like 
anhydrous sodium sulfate, Na 2 S0 4 , decrease in solubility as the 
temperature is raised (134, Fig. 27). If a substance like the 
last named dissolves with evolution of heat, its crystallization 
out of a solution is accompanied by absorption of heat. In 
every case raising the temperature causes that change of solubility 
to occur which involves an absorption of heat. 

The state of chemical equilibrium is shifted in all cases by a 
change of temperature. Now we rind that every chemical reac- 
tion either gives out or absorbs heat. When substances burn, the 
heat given out is very great. In many other reactions the heat 
produced is considerable, while in still others an absorption of 
heat occurs. If a reaction is reversible (all reactions that reach 
a state of equilibrium are, of course, of this class) and produces 



172 . Introduction to General Chemistry 

heat when it goes in one direction, it absorbs an equal amount of 
heat for the same quantity of materials transformed when it goes 
in the opposite direction. 
If the reaction 

H 2 +I 2 ±5 2ffl (264) 

has reached a state of equilibrium at 370 , out of every 1,000 
molecules present 800 will be HI, 100 H 2 , and 100 I 2 . If the 
temperature is then raised to 440 and held constant until a 
new state of equilibrium is reached, the gas mixture will consist 
of 780 molecules of HI, no of H 2 , and no of I 2 . Part of the 
HI has changed to H 2 and I 2 , and the equilibrium may be said 
to have shifted from right to left. At temperatures between 
370 and 440 the change of HI into H 2 and I 2 takes place with 
absorption of heat. We see, then, that raising the temperature 
causes the equilibrium to shift in the direction that involves an 
absorption of heat. Now this is a. perfectly general law for 
chemical changes, just as it is also for physical changes like the 
vaporizing of a liquid and dissolving of a solid. 

289. The Effect of Removing One Product of a Reaction. — 
The reaction represented by the equation 

NaCl-f-H 2 S0 4 ^NaHS0 4 +HCl (103, 253) 

has already been studied rather fully. We may summarize 
the facts briefly, as follows: The action of concentrated sulfuric 
acid on dry salt gives sodium hydrogen sulfate, NaHS0 4 , and 
hydrogen chloride gas, the reaction going nearly to completion 
in the direction of the lower arrow in the equation given above 
if the mixture is warmed. On the other hand, if a cold saturated 
solution of sodium hydrogen sulfate is mixed with concentrated 
hydrochloric acid — that is, a saturated solution of hydrogen 
chloride in water — an abundant precipitate of solid salt, NaCl, is 
formed. This reaction is, we see, just the reverse of the other. 
If now we mix a dilute solution of salt with dilute sulfuric acid, 
we see no visible change. We also see no change upon mixing 
a dilute solution of sodium hydrogen sulfate with dilute hydro- 
chloric acid. 



Chemical Equilibrium 173 

We are now in position to explain all the facts of the foregoing 
paragraph from the standpoint of chemical equilibrium. If 
we bring together dilute solutions of either pair of substances 
in the reaction 

NaCl+H 2 S0 4 ±*NaHS0 4 +HCl, 

the resulting solution probably contains all four substances, side 
by side, in a state of equilibrium. But we cannot notice any 
effect of the mixing, because in the presence of much water all 
four are held completely in solution, since all four are more or 
less readily soluble in water. If, however, but little water is 
present, the least soluble of the four substances, common salt, 
may partially separate. This is the case when a concentrated 
solution of NaHS0 4 is mixed with concentrated HC1. The 
reason is a simple one: the substances taken react partially to 
form NaCl and H 2 S0 4 in the sense of the upper arrow; but the 
amount of NaCl so formed is more than the water present can 
hold in solution; so the excess NaCl separates out in the solid 
form. This separation of NaCl continues* until the four sub- 
stances in the solution have reached amounts which can and do 
exist in equilibrium with one another. Removing the solid 
NaCl which has separated, or adding more solid salt, will in no 
way alter the amounts of any of the four substances contained in 
the solution. 

When concentrated H 2 S0 4 is mixed with dry NaCl, NaHS0 4 
and HC1 begin to be formed. Now HC1 is but slightly soluble 
in concentrated H 2 S0 4 and, being a gas, it at once escapes from 
the mixture. Warming the mixture also promotes the escape 
of the HC1, since the higher the temperature the smaller the 
solubility of the gas in the concentrated H 2 S0 4 . The escape of 
the HC1 gas also has another fundamental effect on the reaction. 
In order that any reaction may reach a state of equilibrium it 
must be reversible; but this reaction cannot go in the reverse 
direction if the HC1 escapes from the reacting mixture as fast 
as it is formed. The result is that if no water is present, con- 
centrated H 2 S0 4 and dry NaCl react practically completely, 
giving solid NaHS0 4 and HC1 gas. 



174 Introduction to General Chemistry 

290. The Action of Steam on Iron and the Reverse Action. — 

When steam is passed over heated iron (29, Fig. 16) hydrogen 
and an oxide of iron are formed. On the other hand, if hydrogen 
is passed over the heated oxide, Fe 3 4 , the products are iron and 
water. The equation for these two reactions, of which one is 
the reverse of the other, is 

3Fe+ 4 H 2 O^Fe 3 4 +4H 2 . 

If we bring together either pair of substances in a closed 
vessel and heat them for some time, a state of equilibrium is 
reached in which all four substances are present. The iron and 
iron oxide are solids, while the water, as steam, is a gas. For 
the condition of equilibrium the relative amounts of steam and 
hydrogen are always the same for a given temperature, no matter 
what proportions of either pair of substances have been used. 
This is the state of affairs if the reaction occurs in a closed vessel. 
But the results are entirely different if the reactions take place 
in a tube holding the solids, through which either steam in the 
one case or hydrogen in the other is passed. If steam is passed 
through a tube containing iron, then the hydrogen which is 
formed is carried along with the excess of steam and has no 
chance to act on the iron oxide which has been formed. There 
is therefore no chance for iron to be formed again, once it has 
been changed to iron oxide. As long as unchanged iron remains 
and the current of steam is continued, the reaction from left to 
right continues. The inevitable result is the complete change 
of the iron to the oxide. On the other hand, if a current of 
hydrogen is passed over heated iron oxide contained in a tube, 
the substances react in the direction from right to left of the 
equation. The steam which is formed passes along with the 
excess of hydrogen, and once having left the tube cannot pos- 
sibly act on the iron to convert it back into oxide, so that this 
change also continues as long as the stream of hydrogen is kept 
up and comes to an end only when all of the iron oxide has been 
reduced to metallic iron. 

291. Conclusions. — We see,, therefore, that a chemical 
reaction like the one just discussed may reach a state of equilib- 



Chemical Equilibrium 175 

rium, if the reverse reaction tends to take place noticeably, and 
if none of the substances involved escape from the vessel in 
which the change takes place ; or it may go to completion in one 
direction or the other if one of the products of either reaction is 
allowed to escape from the scene of action. 

Whether a given reaction reaches a state of equilibrium or 
goes to completion in one direction or the other often depends 
upon the conditions. In the preparation or manufacture of 
chemical substances it is usually very important to cause equi- 
librium reactions to take place as completely as possible in 
order to obtain the maximum possible yields of the desired 
products. 



CHAPTER XIV 
HYDROGEN AND OXYGEN 

292. Hydrogen. — Hydrogen was first recognized in 1766 as 
a distinct substance by Cavendish, a celebrated English chemist, 
who called the gas inflammable air and prepared it by the action 
of acids on metals. It was not until ten years after Cavendish's 
discovery that Lavoisier explained the role played by oxygen 
in combustion and stated the law of the indestructibility of 
matter (21) and thus laid the foundation for the doctrine of the 
elements in its present form. For this reason the classification 
of hydrogen as an element was not possible at the time of its dis- 
covery. In 1 781 Cavendish showed that nothing but water is 
formed when hydrogen burns and thus proved that water is a 
compound of hydrogen and oxygen. The name hydrogen means 
water-former. 

The element occurs in but minute amounts in the free form 
in nature. Water is its most abundant compound ; but it is also 
a constituent of all dry animal and vegetable tissues, forming 
therein principally compounds with carbon, oxygen, and nitrogen. 
Petroleum and natural gas are compounds of hydrogen with 
carbon; coal also contains considerable combined hydrogen. 

293. Preparation of Hydrogen. — We have already learned 
several methods by which free hydrogen can be obtained. These 
may now be briefly reviewed. Hydrogen is formed: 

1. By the electrolysis of water (27), 

2H 2 0->2H 2 +0 2 ; 

2. By the action of water on some metals, as by (a) the burn- 
ing of magnesium wire in steam (28, Fig. 15), . 

Mg+H 2 0->MgO+H 2 , 

(b) the passage of steam over heated iron turnings (29, Fig. 16), 
3 Fe+4H 2 0->Fe 3 4 + 4 H 2 , 

176 



Hydrogen and Oxygen 



177 



(c) the action of sodium or potassium on water (40, 86, Table VI, 
106), 

2 Na+ 2H 2 0->2NaOH+H 2 
2K+2H 2 0-> 2 KOH+H 2 ; 



3, By the action of hydrochloric or sulfuric acid on zinc, 
magnesium, iron, or aluminum, as well as on several other metals, 



Zn-f-2HCl->ZnCl 2 +H 2 
Fe+H 2 S0 4 ^FeS0 4 +H 2 . 



(149) 
(i73) 



294. Making Hydrogen in the Laboratory. — The best labora- 
tory method of making hydrogen consists in treating zinc with 
hydrochloric acid in some form of 
specially constructed gas generator. 
The Kipp apparatus, Fig. 39, is the 
form most extensively used. The 
solution used is made from equal 
volumes of concentrated hydrochloric 
acid and water. The action of this 
generator is very simple in principle. 
Upon opening the stopcock gas 
escapes and allows the acid to rise 
into the middle compartment, where 
it acts upon the zinc and so produces 
a steady flow of hydrogen. When 
the cock is closed the gas formed 
forces the acid downward and causes it to flow from the 
lower into the upper compartment. As soon as the acid is out 
of contact with the zinc all action stops, and no more gas is 
produced until the cock is again opened. The Kipp apparatus 
has one unfortunate defect: since there is but little circulation 
of the solution the acid in contact with the zinc is soon 
exhausted, causing the action to stop while there is still a 
large supply of almost unchanged acid in other parts of the 
apparatus. To start the action again it is necessary to empty 
all the solution and refill with fresh acid; much acid is thus 
wasted. 




Fig. 39 



i 7 8 



Introduction to General Chemistry 



The McCoy apparatus, shown in Fig. 40, has several ad- 
vantages over the Kipp apparatus. The lowest compartment 
is filled as full as possible with granulated or stick zinc, on which 
hydrochloric acid drops at just the rate required to keep up the 
stream of hydrogen that is being drawn from the apparatus. 
When the stopcock is closed the gas which is formed from the 
small excess of acid in the zinc compartment forces the acid from 
the middle to the upper compartment and thus stops the further 

flow of acid upon the zinc. This appa- 
ratus is also conveniently used for 
generating other gases. 

295. The Electrolysis of Water.— 
Compared with metals, pure water is a 
very poor conductor of electricity. The 
addition of a little sulfuric acid increases 
the electrical conductivity of water 
enormously. The sulfuric acid so added 
is not permanently destroyed in the 
course of the electrolysis (27), so that 
very little will serve to promote the 
electrolysis of a large amount of water. 
The exact way in which the acid behaves will be discussed 
later. Ordinarily, poles or electrodes of the elementary metal 
platinum are employed, since most other metals would be 
attacked chemically. The electrode at which the hydrogen is 
liberated is called the negative electrode, or cathode ; the other, 
at which the oxygen appears, is the positive electrode, or 
anode. 

One of the important technical methods of making hydrogen, 
which yields at the same time oxygen, consists in the electrolysis 
of water in which sodium hydroxide is dissolved to make it a 
good conductor. Here the cathode is of iron and the anode of 
carbon. Hydrogen is also obtained in commercial quantities 
as a by-product in the manufacture of caustic soda by the elec- 
trolysis of a solution of common salt. 

Hydrogen is often made for use in balloons by the action of 
dilute sulfuric acid on scrap iron. 




Fig. 40 



Hydrogen and Oxygen 179 

296. The Physical Properties of Hydrogen. — We have already 
learned that hydrogen is colorless ; when perfectly pure it is also 
odorless and tasteless. One liter of the gas at o° and 76 cm. pres- 
sure weighs 0.0899 g-5 an d 22.4 liters, 2 g. approximately. It 
is the lightest of all gases. It can be liquefied, giving a colorless 
liquid which boils at — 253 , or only 20 above absolute zero. At a 
somewhat lower temperature the liquid freezes to a colorless 
solid. 

Hydrogen is but slightly soluble in water: 100 c.c. of water 
dissolves about 2 c.c. of the gas at room temperature. 

The speed of diffusion of hydrogen is greater than that of 
any other gas (191). 

297. The Chemical Properties of Hydrogen. — The most 
important chemical properties of hydrogen have already been 
studied, but may now be briefly reviewed. Hydrogen burns with 
an almost non-luminous flame, which is, however, very much 
hotter than that obtained from ordinary fuel or illuminating gas. 
Water is the product of the reaction. Hydrogen reacts readily 
with hot copper oxide, forming water and copper : 

H 2 +CuO->H 2 0+Cu. (33) 

Hydrogen also acts on other metallic oxides at a red heat, for 
example : 

4H a +Fe 3 4 -»4H a O+3Fe. (290) 

Hydrogen and chlorine, if mixed in equal volumes and ignited, 

or exposed to a bright light, unite with explosive violence, forming 

hydrogen chloride, 

H 2 +C1 2 ->2HC1. (243) 

A jet of burning hydrogen lowered into a jar of chlorine con- 
tinues to burn by reason of the union of the two elements (244) . 
Hydrogen unites with bromine to form hydrogen bromide 
(256) and with iodine to form hydrogen iodide (264). 

298. The Union of Hydrogen and Nitrogen. — A mixture of 
hydrogen and nitrogen does not react at all under ordinary 
conditions. If electric sparks are passed through the mixed 



180 Introduction to General Chemistry 

gases contained in a eudiometer, Fig. 41, a small amount of 
ammonia is formed: 

3 H 2 +N 2 ->2NH 3 . 

The reaction soon reaches a state of equilibrium, because under ' 
the same conditions ammonia is very largely decomposed into 
its elements. As the result of the reverse reaction, a state of 
equilibrium is reached when less than 1 per cent of the ele- 
mentary gases has been converted into ammonia. If the 
■v „ ammonia is absorbed in some suitable way, 

\f^f^ as by union with sulfuric acid, as fast as it is 
formed, then, with continued sparking, the 
formation of ammonia goes on until all 
the hydrogen and nitrogen have united. The 
practical method of making ammonia by this 
reaction will be discussed in chapter xxi. 

299. Heat of Reaction and Flame 
Temperature. — When 1 g. of hydrogen burns, 
about 34,000 calories of heat are produced; 

this is sufficient to heat 340 c.c. of water 
Fig. 41 °^ 

from o° to ioo°. By reason of the great 
amount of heat produced, the flame of hydrogen burning in air 
has a very high temperature. When hydrogen burns in pure 
oxygen instead of in air, the flame is much hotter, but not 
because a greater amount of heat is produced by the burning 
of a given amount of hydrogen, since the quantity of heat is 
the same in the two cases. When hydrogen burns in air, the 
nitrogen, which forms four-fifths by volume of the air, is 
heated to the flame temperature along with the steam formed. 
But in pure oxygen no nitrogen is present, and so the tempera- 
ture reached by the flame is much higher, as there is far less 
material to be heated. 

300. The Oxyhydrogen Blowpipe.— Fig. 42 shows an oxyhy-- 
drogen blowpipe. The two gases mix in the proper proportions 
before issuing from the jet. The temperature of the flame is high 
enough to melt platinum, which cannot be melted in a Bunsen 
flame supplied with fuel or illuminating gas. 




Hydrogen and Oxygen 181 

301. The Limelight. — A very bright light is produced when 
an oxyhydrogen flame strikes a stick of quicklime, by reason of 
the bright white heat to which the lime is raised. This is the 
so-called limelight, which was very extensively used before the 
electric arc light was perfected and which is still frequently 
used in rural communities. In place of lime other difficultly 
fusible white oxides may be employed. For this purpose 
thorium oxide containing 1 per cent of cerium oxide is much 
superior to lime. 

302. Ignition Temperature. — A mixture of hydrogen and 
oxygen in their combining proportion remains unchanged for any 




Fig. 42 

length of time at room temperature, but if brought in contact 
with a flame or electric spark it explodes violently. The explo- 
sion is due to the great increase in volume of the reaction product, 
steam, caused by the almost instantaneous reaction, with its 
attendant heat production. But we may inquire why a reaction 
which does not take place at room temperature can become explo- 
sive. Investigation shows that a mixture of the two gases reacts 
perceptibly at 450 , and that the formation of water goes on 
faster the higher the temperature, but that the mixture does not 
become explosive until the temperature reaches about 6oo°. It 
is easy to see why explosion finally occurs when the temperature 
is raised. While the reaction is taking place slowly, heat is being 
produced by the chemical change; below 6oo° the rate of change 
is so slow that heat is lost by the gas mixture faster than it is 
produced. Above 6oo° the reaction goes faster, so that heat is 
produced more rapidly than it is lost, and this causes the gas 
mixture to grow hotter; and the hotter it gets the faster the 
reaction goes, until soon it proceeds with enormous rapidity, and 
this constitutes an explosion. For any combustible substance 
there is some temperature to which it must be heated before its 



1 82 Introduction to General Chemistry 

rate of production of heat by union with oxygen exceeds its rate 
of loss of heat; if heated to this temperature the substance takes 
fire and continues to burn. This point is called the ignition 
temperature. 

Hydrogen, issuing from a jet, burns quietly in air when 
ignited. This is because the actual union with oxygen can occur 
only as fast as the two gases can reach one another by diffusion 
(191), one from the jet, the other from the surrounding air. 
The flame is the reacting gas mixture, which is raised to incan- 
descence by the great heat produced by the union. 

303. Platinum as a Catalytic Agent. — The element platinum 
can be deposited on asbestos as a thin, spongy coating by dipping 
a bit of fibrous asbestos in a solution of platinum chloride, 
PtCl 4 , drying the material and holding it in a Bunsen flame for 
a minute. The salt decomposes into its elements thus: 

PtCl 4 ^Pt+2Cl 2 . 

If a jet of cold hydrogen gas is directed on the cold platinized 
asbestos, the latter gets red-hot and sets fire to the hydrogen. 
Spongy platinum absorbs gases to a marked extent. Heat is 
produced in this way, and this ignites the intimate mixture of 
hydrogen and oxygen condensed on the surface of the metal. 
The platinum itself is entirely unchanged and will continue 
active in this way indefinitely. A substance which initiates or 
promotes a chemical reaction without itself being changed is 
called a catalytic agent (239). 

304. The Use of Hydrogen in Balloons. — The lifting power 
of a balloon filled with hydrogen can easily be calculated. Since 
1 liter of air under standard conditions weighs 1 . 29 g. and 
1 liter of hydrogen weighs 0.09 g., the difference, 1. 2 g., repre- 
sents the lifting power per liter of capacity of a balloon. At 
higher temperature and lower pressure the lifting power is 
smaller. If a Zeppelin has a capacity of 5,000,000 liters its 
lifting power will be about 6,000 kilos, or more than 13,000 
pounds. 

305. Oxygen. — Oxygen in the form of compounds makes up 
about one-half by weight of the matter forming the crust of the 



Hydrogen and Oxygen 183 

earth. It also constitutes 89 per cent by weight of water and 
21 per cent by volume of the air. The oxygen of the air is not 
chemically combined bu.t is only mixed with nitrogen and small 
amounts of other gases present. Oxygen was first prepared by 
Priestley, in England (1774), by heating mercuric oxide, 

2 HgO->2Hg+0 2 . 

At practically the same time Scheele, in Sweden, made oxygen 
by this method and also by heating potassium nitrate, 

2KN0 3 -> 2 KN0 2 +0 2 . 

The salt KN0 2 is called potassium nitrite. The name oxygen, 
which means acid-former, was given to the gas by Lavoisier, who 
believed that it was a necessary constituent of all acids. At 
that time hydrochloric acid, which was called muriatic or marine 
acid, was thought to contain oxygen. We now know many 
other acids which do not contain oxygen. 

306. The Preparation of Oxygen. — We have already learned 
several ways by which oxygen may be made. The heating of 
mercuric oxide and the electrolysis of water (14, 295) have already 
been fully studied. We have also seen (245) that oxygen is 
formed when chlorine water is exposed to sunlight: 

2Cl 2 +2H 2 0->0 2 +4HCl. 

The heating of certain salts which are rich in oxygen is also 
a simple way of making the gas. The behavior of potassium 
nitrate, KN0 3 , is given in the preceding paragraph. Potassium 
chlorate, KC10 3 , is easily decomposed by heat according to the 
following equation: 

2KC10 3 ->2KCH-30 2 . 

This last reaction, is the one usually employed in making small 
amounts of oxygen in the laboratory. It can be carried out in 
a test tube, a small flask, or a retort. The crystals of potassium 
chlorate first melt, and at a little higher temperature the liquid 
seems to boil, by reason of the oxygen given off. 



184 Introduction to General Chemistry 

The change of the chlorate, KC10 3 , into chloride, KC1, does 
not take place completely in one step. The first stage of the 
reaction is probably represented by the # equation, 

ioKC10 3 -> 4 KCl+6KC10 4 +30 2 . 

The salt KC10 4 , called potassium perchlorate, can also be decom- 
posed by heat, thus : 

KC10 4 ^KCl+20 2 . 

This last reaction requires a higher temperature than the first. 
If the heating of the chlorate is stopped when about one-fifth of 
its total oxygen has been given off, KC10 4 will be found in the 
residue. 

In making oxygen from potassium chlorate two precautions 
should be observed : first, the material must be entirely free from 
bits of wood, paper, etc., which are easily combustible; and, sec- 
ondly, the heating must be gentle, as otherwise the decomposition 
may occur explosively. 

When powdered potassium chlorate is mixed with about half 
its weight of manganese dioxide, Mn0 2 , it will give off its 
oxygen rapidly at a temperature far below that at which the 
pure chlorate starts to decompose. Since manganese dioxide 
alone does not give off any of its oxygen until a rather high 
temperature is reached, and is not changed itself in promoting 
the decomposition of the potassium chlorate, we must consider 
that the former substance acts only as a catalytic agent in pro- 
moting the decomposition of potassium chlorate. 

307. Oxygen from Sodium Peroxide. — Sodium peroxide, 
Na 2 2 , is a solid made by burning metallic sodium, 

2Na+0 2 ^Na 2 2 . 

The trade name of the material is oxone; it is supplied in the 
form of lumps or sticks. Water acts on it as follows: 

2Na 2 2 + 2H 2 0-> 4 NaOH+0 2 . 

By dropping water on lumps of oxone contained in a suitable 
apparatus (236, Fig. 30) a steady stream of oxygen is obtained. 



i Hydrogen and Oxygen 185 

The method is rather expensive, but it is very convenient, since 
the action stops when the supply of water is turned off and can 
be started again at will. 

308. Oxygen from Other Oxides. — Lead dioxide, Pb0 2 , when 

strongly heated gives oxygen and lead monoxide, or litharge, 

PbO: 

2Pb0 2 ->2PbO+0 2 . 

Manganese dioxide is also decomposed at a high temperature, 

thus: 

3 Mn0 2 ->Mn 3 4 +0 2 . 

309. Technical Methods of Making Oxygen. — The elec- 
trolysis of water is an important technical method of making 
oxygen. It also yields hydrogen and has already been described. 
By far the larger part of the oxygen of commerce is made from 
liquid air. This substance is a mixture of liquid oxygen and 
liquid nitrogen. The latter boils about n° lower than the 
former, whose boiling-point is — 183 , and therefore distils off 
first when liquid air is allowed to evaporate, leaving nearly pure 
oxygen. This is stored under pressure in steel tanks and brought 
on the market. 

310. Brin's Process. — Brin's process, formerly used tech- 
nically, is a method of obtaining oxygen from the air by means of 
barium oxide, BaO. This oxide unites with more oxygen at a 
red heat, forming barium peroxide : 

2BaO+0 2 ^2Ba0 2 . 

This is a reversible reaction, which at a constant temperature 
will go in one direction or the other with change of pressure. 
In practice, air is pumped under pressure into a vessel contain- 
ing BaO at 700 , and when all the oxide has been changed into 
Ba0 2 the nitrogen present is allowed to escape. By reducing 
the pressure with a vacuum pump the Ba0 2 is caused to decom- 
pose completely, yielding nearly pure oxygen. 

311. Oxygen from Plants. — Growing plants absorb carbon 
dioxide from the air. They also take up water through their 
roots. In some manner, not fully understood, carbon dioxide 
and water react under the influence of sunlight to form such 



1 86 Introduction to General Chemistry 

principal plant constituents as starch, cellulose, and sugar, 
together with oxygen, which is given off to the air. The per- 
centage of oxygen in the air would soon decrease if it were not 
maintained by growing plants. 

312. The Physical Properties of Oxygen. — It is, of course, 
obvious that oxygen is colorless, odorless, and tasteless. One 
liter weighs 1.429 g. and 22.4 liters about 32 g., corresponding 
to the formula 2 . Liquid oxygen is pale blue in color; it 
boils at — 183 . At o°, 100 c.c. of water dissolves about 5 c.c. 
of oxygen; at 20 , about 3 c.c. (125). 

313. The Chemical Properties of Oxygen. — We have already 
learned that combustion was first explained by Lavoisier in 
1774 as due to union of the burning substance with the oxygen 
of the air (13-15). All the elements so far studied, except 
fluorine, form oxides. This does not mean that all these ele- 
ments burn, since some oxides, like those of chlorine and silver, 
can only be made indirectly (172). The oxides of metallic 
elements, by union with water, form hydroxides which are 
bases, for example: 

CaO+H 2 0->Ca(OH) 2 . 

The oxides of non-metallic elements, including carbon, sulfur, 
nitrogen, phosphorus, and the halogens (except flourine), give, 
with water, acids. The following equations will serve as illus- 
trations of such reactions, some of which have already been 
studied; the others will be studied later. 

C0 2 +H 2 0->H 2 C0 3 , carbonic acid, 

S0 3 +H 2 0->H 2 S0 4 , sulfuric acid, 
N 2 5 -fH 2 0->2HN0 3 , nitric acid, 
P 2 5 -|-3H 2 0->2H 3 P04, phosphoric acid, 

I 2 S +H 2 0->2HI0 3 , iodic acid. 

An oxide which by union with water forms an acid is often 
called the anhydride of the acid. 

314. Respiration. — Animals breathe air in order to obtain 
oxygen. The blood contains a complex substance, haemo- 
globin, which forms with oxygen a compound, oxyhaemoglobin, 
which easily decomposes reversibly into oxygen and haemo- 



Hydrogen and Oxygen 187 

globin. This- is a typical equilibrium reaction: when oxygen, 
at the pressure at which it exists in the air, comes in contact with 
the blood in the lungs the compound is formed, 1 g. of haemo- 
globin uniting with 1.3 c.c. of oxygen; when the blood reaches 
the tissues, which take up oxygen, the compound decomposes 
and the haemoglobin is carried by the blood bapk to the lungs, 
where it again takes up fresh oxygen from the air. 

315. Uses of Oxygen. The Oxyacetylene Torch. — The use 
of oxygen in the oxyhydrogen blowpipe has already been men- 
tioned. By substituting acetylene for hydrogen in a blowpipe 
similarly constructed we get the oxyacetylene torch, which 
gives an intensely hot flame. It is extensively used for welding 
and for cutting iron and steel. 

Oxygen is used in several analytical processes, such as those 
studied in chapter iv. 

Deposits of carbon in the cylinders of gasoline engines are 
often removed by burning out with oxygen. Since iron burns 
also rather readily in a stream of oxygen, care must be taken to 
avoid injuring the cylinder in this way. 

316. Ozone. — When a silent electric discharge passes through 
oxygen a very remarkable change is produced ; there is a decrease 
in volume, and a gas having a powerful irritating odor is pro- 
duced. The new gas is ozone. The simplest form of apparatus 



Wffll 




Fig. 43 

used for making ozone is shown in Fig. 43 . It is a double-walled 
glass tube having the outside of the outer tube and the inside of 
the inner tube coated with tin foil. These coatings are con- 
nected by wires to the terminals of an induction coil. When the 
coil is set in action and a slow stream of oxygen is passed through 
the space between the outer and the inner tubes, the issuing gas 
is found to contain ozone. The peculiar odor of the air in the 



1 88 Introduction to General Chemistry 

neighborhood of powerful electrical machinery is due to ozone. 
Ozone is very much more active as an oxidizing agent than 
oxygen. Mercury shaken with ozone is very quickly oxidized. 
Ozone also sets iodine free from a solution of an iodide. 

If nothing but oxygen is needed to produce ozone — and such 
is actually the case — what then is the cause of the remarkable 
change in properties? In the first place it was noticed that a 
decrease in volume occurs when ozone is formed from oxygen. 
On the other hand, when ozone is changed to oxygen, as may be 
done by heating the former, the volume of the oxygen, when it 
is again cooled to the original temperature, is greater than that 
of the ozone. In fact, three volumes of oxygen give exactly two 
of ozone and vice versa. The density of ozone is one-half 
greater than that of oxygen. While 22.4 liters of oxygen weigh 
32 g., the same volume of ozone weighs 48 g. Ozone is oxygen 
in another form. If for oxygen we write the formula 2 , we must 
write 3 as the formula of ozone. The molecules of ozone 
differ from those of oxygen by containing three instead of two 
atoms. We may write the reversible equation for the relation 
between oxygen and ozone thus: 

• 30 2 ^20 3 . 

When ozone acts on mercury, for example, the action is as follows : 

Hg+0 3 ^HgO+0 2 . 

Only one-third of the oxygen of ozone is active, the balance changing 
into ordinary oxygen. Iodides are oxidized by ozone, thus: 

2KI+0 3 +H 2 0^ 2 K0H+I 2 +0 2 . 

The liberated iodine may be recognized by its action on starch. 
Minute amounts of ozone may be recognized in this way, 
although the test is not conclusive proof of the presence of ozone, 
since many other substances also set iodine free from iodides. 

317. Ozone as a Germicide. — Since ozone is a very powerful 
ozidizing agent, it is not surprising that it should readily destroy 
germs. It has been found that impure water containing even 
1,000,000 bacteria per c.c. is completely sterilized by intimate 



Hydrogen and Oxygen 189 

contact with an equal volume of air containing 2 g. of ozone 
per cubic meter. In a number of important cities of Europe the 
entire municipal water supply is purified by means of ozone. 
Disinfection by chlorine is more popular. 

318. Hydrogen Peroxide, H 2 2 . — The well-known household 

antiseptic and disinfectant, hydrogen peroxide, H 2 2 , is a 3 per 

cent solution of this substance in water. Some hydrogen 

peroxide is formed by the action of sodium peroxide on ice 

water, thus: 

Na 2 2 + 2 H 2 ±5 H 2 2 + 2 NaOH. 

If water is dropped on sodium peroxide the material becomes 
very hot, and oxygen and water instead of hydrogen peroxide 
are formed (307). This is because the latter substance easily 
decomposes if hot, especially in the presence of caustic soda: 

2H 2 2 ^2H 2 0+0 2 . 

319. Preparation of Hydrogen Peroxide. — Barium peroxide, 
Ba0 2 , the formation of which was discussed in connection with 
Brin's process of making oxygen (310), reacts with dilute sulfuric 
acid to form hydrogen peroxide and barium sulfate : 

Ba0 2 +H 2 S0 4 ±5H 2 2 +BaS0 4 . 

Since barium sulfate is insoluble in water a very pure solu- 
tion of hydrogen peroxide is easily obtained. The reaction is 
best carried out by adding finely powdered barium peroxide, 
suspended in water, very gradually to ice-cold, diluted sul- 
furic acid. The precipitate of barium sulfate is allowed to 
settle, leaving a clear solution of hydrogen peroxide. This 
solution must be made as nearly neutral as possible, other- 
wise it will decompose more or less rapidly into oxygen and 
water. 

By cautious evaporation, at a moderate temperature in a 
partial vacuum, a dilute solution of hydrogen peroxide may be 
freed from most of its water; the resulting concentrated solution 
when cooled to — io° deposits crystals of H 2 2 . 

320. Properties of Hydrogen Peroxide. — At ordinary tem- 
peratures pure hydrogen peroxide is a colorless liquid which will 



190 Introduction to General Chemistry 

mix with water in all proportions. It freezes at — 2 . It does 
not boil without decomposition, and when strongly heated it is 
liable to explode, water and oxygen being the products. The 
speed of decomposition of hydrogen peroxide at ordinary 
temperatures is' greatly influenced by the presence of other sub- 
stances which act as catalytic agents. Finely divided metals 
like platinum and gold cause hydrogen peroxide to decompose 
rapidly. Manganese dioxide behaves similarly. In these reac- 
tions neither the metals nor the manganese dioxide are changed. 
They are catalytic agents (303). 

Hydrogen peroxide seems to have the property of an acid, 
since it combines with some bases to form compounds which 
may be considered salts. For example, with barium hydroxide, 
Ba(OH) 2 , it reacts thus: 

H 2 2 +Ba(OH) 2 + 6H 2 O^Ba0 2 - 8H 2 0. 

The product consists of white crystals, rather difficultly soluble 
in water. 

Hydrogen peroxide is often used as a bleaching agent for 
plant and animal substances, such as hair, feathers, silk, ivory, 
and straw. 

The most characteristic property of hydrogen peroxide is its 
great tendency to give up oxygen and thus to act on substances 
capable of reacting with oxygen. For example, with hydriodic 
acid it gives iodine and water: 

H 2 2 +2HI^I 2 + 2 H 2 0. 

The free iodine can easily be recognized by the blue color which 
it gives with a starch solution. Instead of using hydriodic 
acid we may use a solution of potassium or sodium iodide to 
which hydrochloric acid has been added, since the solution will 
then contain some hydriodic acid formed as follows: 

KI+HC1^KCI+HI. 

321. Detection of Hydrogen Peroxide.— A very delicate 
reaction which serves to detect small quantities of hydrogen 



Hydrogen and Oxygen 191 

peroxide is that which occurs when a solution of this substance 
is mixed with a little sulfuric acid and a very dilute solution of 
potassium dichromate, K 2 Cr 2 7 . The latter substance contains 
the element chromium, Cr, as one of its constituents. The 
solution turns blue, and when a little, ether is added and shaken 
up with the blue solution the ether dissolves the blue substance. 
If the mixture is allowed to stand a minute or two the blue ether 
solution separates from the water solution, on which it floats as a 
blue layer. 

Other reactions of hydrogen peroxide are discussed in the 
following chapter (347, 348). 

322. Peroxides and Dioxides. — We have just learned that 
hydrogen peroxide is formed by the action of dilute acids on 
Na 2 2 and Ba0 2 , and we might therefore be inclined to expect 
that we should also get H 2 2 by the action of acids on Pb0 2 and 
Mn0 2 . But this is not the case; no H 2 2 can be obtained in 
any way from these last-mentioned oxides. For this reason 
these oxides of lead and manganese are called dioxides to dis- 
tinguish them as a class from those which yield H 2 2 and which 
are called peroxides. Thus we call Ba0 2 barium peroxide and 
Pb0 2 lead dioxide. 

323. Graphic Formulae.; — We have so far considered that 
oxygen has a valence (183) of two, or is bivalent, since in water 
two symbol weights of hydrogen are united with one of oxygen. 
But what then is the valence of oxygen in H 2 2 ? In order to be 
able to answer this question, we must consider the matter of 
valence from the standpoint of the atomic-molecular hypothesis. 
We have learned (221) that the molecule of water is made up 
of two atoms of hydrogen and one of oxygen. Since all the 
molecules of water are made up in just this fashion, it would seem 
to follow that the three atoms must be related to one another 
in some very definite way. We may think of them as being 
joined to one another, in which case there are the two possibilities 
indicated by the following graphic formulae in which the lines 
joining the symbols are called bonds. 

, (1) H-H-0 (2) H-O-H. 



192 Introduction to General Chemistry 

The first graphic formula indicates that one of the hydrogen 
atoms is attached on the one hand to the atom of oxygen and 
on the other to the second atom of hydrogen. Formula (2) indi- 
cates that it is the oxygen atom which is attached on either hand 
to an atom hydrogen. It is obvious that the second formula 
is the more consistent, since in it both atoms of hydrogen are 
attached by single bonds to the atom of oxygen which holds an 
atom of hydrogen by each of its two bonds. In formula (1) the 
middle hydrogen atom is represented as having two bonds, while 
the other hydrogen atom and also the atom of oxygen are shown 
as having but one bond each. Since we think of all atoms of 
hydrogen as being alike, we must reject the first formula in favor 
of the second. 

When viewed in the above-mentioned manner the valence of 
an element is seen to be the holding capacity of its atoms for atoms 
of hydrogen or other univalent elements like chlorine. We may 
therefore think of an atom of oxygen which is bivalent as having 
two valence bonds, each of which can hold one atom of a univalent 
element. 

324. The Graphic Formulae of Peroxides. — We are now 
prepared to consider the question of the valence of oxygen in 
H 2 2 and the graphic formula of this substance. At the outset 
it may be stated that the valence of hydrogen is considered by 
chemists to be invariably one. If the valence of oxygen is taken 
to be two, there is but one possible graphic formula, namely: 

H-O-O-H. 

This is the commonly accepted formula. For sodium peroxide 
we then have the formula, 

Na-O-O-Na, 
and for barium peroxide, 

X 



Hydrogen and Oxygen 193 

In their dioxides, manganese and lead have without doubt 
a valence of four, or are tetravalent, and oxygen is, as usual, 
bivalent. We therefore write for these oxides the formulae 

= Mn = and = Pb = 0, 

thereby indicating that each oxygen atom is attached to an 
atom of manganese or lead by two bonds, or in other words by a 
double bond. 

The monoxides of manganese, MnO, and lead, PbO, have 
their atoms doubly bound, thus : 

Mn=OandPb = 0. 
In these oxides both metals are bivalent. 



CHAPTER XV 
OXIDATION AND REDUCTION 

325. Oxidation. — When a substance unites with oxygen it 
is said to be oxidized. Hydrogen when burned is oxidized, giv- 
ing water. Metals are said to be oxidized when they combine 
with oxygen; for example, 

2Cu+0 2 ->2CuO. 

By certain indirect methods a lower oxide of copper, cuprous 
oxide, Cu 2 0, can be made. This oxide can unite with more 
oxygen if heated in air or in oxygen and form the common oxide, 
CuO, which is known also as cupric oxide, in order to distinguish 
it from the lower oxide: 

2Cu 2 0+0 2 ->4CuO. 

We say in this case that cuprous oxide has been oxidized to cupric 
oxide. 

The action of oxygen gas on hydrogen chloride at a high 
temperature (239) proceeds according the equation 

4 HCl+0 2 ->2Cl 2 +2H 2 0, 

and in consequence we say that the hydrogen chloride has been 
oxidized. 

326. Oxidizing Agents. — Very often substances may be 
oxidized by compounds of oxygen as well as by oxygen itself. 
For example, heated copper oxide, CuO, oxidizes hydrogen and 
all its compounds, such as ammonia, acetylene, etc. We say, 
therefore, that copper oxide is an oxidizing agent. Any sub- 
stance which oxidizes another is called an oxidizing agent. Lead 
dioxide and manganese dioxide are powerful oxidizing agents, 
as shown by the fact that each is able to oxidize hydrochloric 

acid, 

Mn0 2 + 4 HCl->MnCl 2 +Cl 2 +2H 2 0. (234) 

194 



Oxidation and Reduction 195 

Potassium permanganate, KMn0 4 , which also easily oxi- 
dizes hydrochloric acid,, is one of the most powerful of all oxidiz- 
ing agents. It acts according to the equation 

2KMn0 4 + i6HCl^ 2 KCl+ 2 MnCl 2 + 5CL+ 8H 2 0. (235) 

Nitric acid and its salts (104), the nitrates, are good oxidiz- 
ing agents. Gunpowder is a mixture of finely powdered potas- 
sium nitrate, charcoal, and sulfur. The explosion of gunpowder 
is due to the extremely rapid oxidation of the charcoal (carbon) 
and sulfur to carbon dioxide, C0 2 , and sulfur dioxide, S0 2 , the 
oxygen being furnished by the potassium nitrate, KN0 3 . The 
other products of the explosion are nitrogen and potassium sul- 
fide, K 2 S. 

Potassium chlorate, KC10 3 , is a powerful oxidizing agent, 
readily giving up all its oxygen to oxidizable substances and 
leaving the chloride, KC1. 

Sulfuric acid (93) is capable of oxidizing some substances; 
hot sulfuric acid acts on charcoal thus: 

2H 2 S0 4 +C->C0 2 +2S0 2 +2H 2 0. 

It is probable that each molecule of sulfuric acid first loses one 
atom of oxygen giving a molecule of sulfurous acid, H 2 S0 3 , and 
that this substance, which is unstable, then decomposes into 
sulfur dioxide and water: 

H 2 S0 3 ->S0 2 +H 2 0. (286) 

327. Reduction and Reducing Agents. — When hydrogen is 
oxidized by hot copper oxide, thus: 

CuO+H 2 ->Cu+H 2 0, 

the copper oxide is said to be reduced to metallic copper. In 
consequence we call hydrogen a reducing agent. Any substance, 
as, for example, acetylene or methane, which can reduce copper 
oxide is also called a reducing agent. In any reaction, if one 
substance is oxidized the oxidizing agent is by necessity reduced; 
oxidation and reduction always go on together. All substances 
which are acted upon by oxidizing agents are, of 'course, redu- 
cing agents. 



196 Introduction to General Chemistry 

328. Carbon as a Reducing Agent. — Since charcoal, which 
is nearly pure carbon, burns readily, it Js capable of taking up 
oxygen from oxidizing agents and is therefore a good reducing 
agent. A mixture of powdered copper oxide and charcoal reacts 
vigorously, if strongly heated, giving copper and carbon dioxide: 

2CuO+C->2Cu+C0 2 . 

In this reaction the copper oxide is the oxidizing agent and the 
charcoal (carbon) the reducing agent. 

Many other metallic oxides can be reduced in a similar man- 
ner by carbon. In place of charcoal, coke or coal, which are 
largely carbon, may be used. Thus ferric oxide, Fe 2 3 , which 
in the form of the mineral hematite is the most important ore 
of iron, is reduced by coke at a white heat to metallic iron. The 
result may be represented by the equation 

2Fe 2 3 +3C->4Fe+ 3 C0 2 , 

although it is very probable that the reaction is less simple under 
the conditions actually met with in practice. 

329. Carbon Monoxide as a Reducing Agent. — When carbon 
is burned in a deficient supply of air, carbon monoxide, CO, is 
formed instead of dioxide: 

2C + 2 ->2CO. 

v 
This is a colorless, odorless, and very poisonous gas which will 

burn with a nearly non-luminous flame to form carbon dioxide, 

2CO+0 2 ->2C0 2 . 

Some oxidizing agents are able to oxidize carbon only to 
monoxide and not to dioxide. Zinc oxide behaves in this way: 

ZnO+C->Zn+CO. 

This is the reaction by which zinc is made from its ores. The 

reaction between ferric oxide and carbon can also give carbon 

monoxide, 

Fe 2 3 +3C->2Fe+3CO. 

But carbon monoxide can also reduce ferric oxide: 
Fe 2 3 +3CO->2Fe+3C0 2 . 



Oxidation and Reduction 197 

The last two equations doubtless represent the steps by which 
ferric oxide and carbon react to give iron and carbon dioxide 
(328). 

330. Aluminum as a Reducing Agent. — Metallic aluminum 
unites vigorously with oxygen at a white heat, although it has 
no tendency to oxidize in the air at ordinary temperatures. The 
burning of aluminum occurs thus : 

4Al+30 2 ->2Al 2 3 . 

When a mixture of powdered aluminum and ferric oxide is 
strongly heated a very violent reaction takes place, giving iron 
and aluminum oxide: 

2Al+FeAr>2Fe+Al 2 3 - 

The mixture of aluminum and ferric oxide has been given the 
trade name of thermite by its inventor, Goldschmidt, who uses 
it to make small quantities of molten iron for the repair of broken 
iron castings, etc. 

Many other metallic oxides can also be reduced by aluminum. 

331. Oxidation Considered as a Change of Valence. — We 
have already learned (173) that iron forms two series of com- 
pounds , ferrous and ferric, as illustrated by the following formulae: 



Ferrous Compounds 


Ferric Compounds 


FeO 


Fe 2 3 


Fe(OH) 2 


Fe(OH) 3 


FeCl 2 


FeCl 3 


FeBr 2 


FeBr 3 


Fe(N0 3 ) 2 


Fe(N0 3 ) 3 


FeS0 4 


Fe 2 (S0 4 ) 3 



The valence of iron is two in ferrous compounds and three in 
ferric. According to this usage of the term valence we should 
be forced to say that the valence of free or uncombined iron is 
zero. 

If free iron is changed into ferrous oxide, 

2Fe+0 2 ->2FeO, 



198 Introduction to General Chemistry 

it is oxidized, and its valence is increased from zero to two. 
Moreover, if ferrous oxide is changed into ferric oxide, 

4FeO+0 2 ->2Fe 2 3 , 

it is also plain that the iron is further oxidized, and that its 
valence has increased from two to three. It is customary to 
say that ferric oxide is a higher oxide of iron than ferrous oxide ; 
or that iron in ferric oxide is in a higher state of oxidation than 
in ferrous oxide. 

In the case of iron and its oxides we see that the oxidation of 
iron and the increase in its valence go hand in hand. 

With respect to other elements that unite with oxygen we 
also find that their oxidation results in an increase in their 
valence. A few additional examples will help to illustrate this 
point. In the change of copper into cuprous oxide, Cu 2 (325), 
the oxidation of the copper is accompanied by an increase of its 
valence from zero to one. In cuprous oxide copper is univalent 
(146). In the oxidation of cuprous oxide to cupric oxide, 

2Cu 2 0+0 2 -> 4 CuO, (325) 

the' valence of copper is increased from one to two. In cupric 
oxide copper is bivalent (146). 

When carbon is oxidized to carbon monoxide, 

2C+0 2 ->2CO, (329) 

the valence of Carbon is increased from zero to two (carbon is 
bivalent in carbon monoxide). In the oxidation of carbon 
monoxide to carbon dioxide, 

2CO+0 2 ->2C0 2 , (329) 

the valence of carbon is increased to four, carbon becoming 
quadrivalent. 

332. A Broader Meaning of the Term Oxidation. — Since in 
the change of any ferrous compound into the corresponding 
ferric compound (173, 331) the- valence of iron always increases 
from two to three, all such changes may well be considered to 
be of the same class. It has become the custom among chemists 



Oxidation and Reduction 199 

to call such increase of valence of iron an oxidation of the iron 
irrespective of the nature of the element or radical combined 
with the iron. Thus in the reaction 

2FeCl 2 +Cl 2 ->2FeCl 3 , 

whereby ferrous chloride is changed to ferric chloride, we say 
that the iron has been oxidized. The only question that should 
arise here is : Why call this increase in valence of iron an oxida- 
tion in cases where no oxygen is involved? We can only say 
that it is a custom sanctioned by long and universal usage. 

By way of further illustration of the use of the term oxida- 
tion in its broader sense we may cite the following examples. 
When metallic sodium is changed into chloride, NaCl, or nitrate, 
NaN0 3 , its valence is increased from zero to one, and we say 
that the sodium has been oxidized. When zinc is changed into 
oxide, ZnO; sulfate, ZnS0 4 ; chloride. ZnCl 2 ; or nitrate, Zn(N0 3 ) 2 
(148), we say that the zinc has undergone oxidation; and further- 
more, since in all these compounds zinc is bivalent, we say that 
zinc in all these compounds is in the same state or stage of oxida- 
tion. In fact, zinc in its compounds is always bivalent. 

333. Review of Other Elements with Variable Valence. — 
Iron is not the only element having a variable valence. We 
have already seen (179, 180) that mercury also forms two series 
of compounds, the mercurous, in which the element has a valence 
of one, and the mercuric, where the valence is two, as illustrated 
by the following formulae: 



rous Compounds 


Mercuric Compour 


Hg 3 


HgO 


HgCl 


HgCl, 


Hgl 


Hgl, 


HgN0 3 


Hg(N0 3 ) 2 


Hg 2 S0 4 


HgS0 4 



Mercurous compounds are converted into mercuric by oxida- 
tion, and mercuric into mercurous by reduction. 

Copper also forms two series of compounds, cuprous and 
cupric. W T e know cuprous oxide, Cu 2 0, and cuprous chloride, 
CuCl, as well as the commoner cupric compounds, such as cupric 



200 Introduction to General Chemistry 

oxide, CuO, cupric chloride, CuCl 2 , cupric sulfate, CuS0 4 , 
etc. (165). 

It will be noted that compounds representing the lower state 
of oxidation have names ending in ous, while those corresponding 
to the higher state of oxidation end in ic. 

334. Another Class of Oxidation Reactions. — We have 
given as one illustration of an oxidation reaction the action of 
manganese dioxide on hydrochloric acid : 

Mn0 2 + 4 HCl->MnCl 2 +Cl 2 +2H 2 (326) 

In this reaction the manganese dioxide is the oxidizing agent. 
The chlorine of the hydrochloric acid in being set free has its 
valence decreased from one to zero. The valence of the hydro- 
gen remains unchanged in the reaction: it is neither oxidized 
nor reduced. Now we see that when combined chlorine, as in a 
chloride, is set free its valence is decreased, and that in this change 
the chlorine is oxidized. In the case of metallic elements oxida- 
tion involves increase of valence (331). We see, therefore, that 
in the case of chlorine, a non-metallic element, its oxidation in 
cases like that cited involves a decrease in its valence. Other 
non-metallic elements, like bromine, iodine, and sulfur, when set 
free by oxidizing agents from their compounds with hydrogen 
or metals are oxidized, while at the same time there occurs a 
decrease in valence. 

335. Reduction and Change of Valence. — That valence 
changes accompany reduction, as well as oxidation, will be at 
once apparent by the consideration of any reaction in which 
reduction takes place. Take, for example, the simple case of 
the reduction of cupric oxide by hydrogen, 

CuO+H 2 ->Cu+H 2 0. (327) 

The copper is reduced to the free state, its valence changing 
from two to zero, while the oxygen merely changes partners 
without change of valence and is therefore neither oxidized nor 
reduced. When any oxide of a metal is reduced by hydrogen, 
carbon, or any other reducing agent, the valence of the metal is also 
reduced or lowered. 



Oxidation and Reduction 201 

On the other hand, in the reaction 
C1 2 +H 2 -> 2 HC1 

we see that the reduction of the non-metallic element chlorine 
is accompanied by an increase in its valence from zero to one. 
In general, when a non-metallic element like a halogen, sulfur, or 
oxygen itself unites with hydrogen or a metal, the non-metal is 
reduced, while concurrently its valence is increased. 

In an earlier paragraph (330) we learned that aluminum can 
act as a reducing agent, as indicated by the following equation: 
2Al+Fe 2 3 ^2Fe+Al 2 3 - 

Aluminum also reacts very readily with a solution of ferric 
chloride to reduce it to ferrous chloride: 

3FeCl 3 +A1^3FeCl 2 +AlCl 3 . 

Other metals, like zinc and magnesium, are also good reducing 
agents, as indicated by the following equations 

Fe 2 (S0 4 ) 3 +Zn-> 2 FeS0 4 +ZnS0 4 , 
2FeBr 3 +Mg->2FeBr 2 +MgBr 2 . 

336. Discussion. — It is not profitable at this time to go more 
deeply into generalizations regarding the valence changes accom- 
panying oxidation and reduction, for the reason that this subject 
involves other matters which must be taken up later (chap, xx), 
and without a knowledge of which no satisfactory or complete 
discussion of this important subject is possible. However, we 
have gone far enough to see that in every oxidation and reduction 
the valencies of the elements oxidized and reduced change in a 
systematic fashion. Any reagent that can bring about such 
changes of valence is either an oxidizing or a reducing agent, 
even though it does not contain either oxygen or hydrogen. 
Thus, for example, in the reaction 

2FeCl 2 +Cl 2 ^2FeCl 3 , (332) 

in which the iron of the ferrous chloride is oxidized to the ferric 
state, the free . chlorine used is the oxidizing agent. We call 
ferrous chloride the reducing agent, although, to be exact, it is 
only the iron of this salt that has any reducing action, since the 



202 Introduction to General Chemistry 

chlorine already in FeCl 2 is neither oxidized nor reduced in the 
change to FeCl 3 . 

337. Two Important Kinds of Reactions. — But very few of 
the reactions of substances in solution studied in chapters prior 
to the present involve oxidation and reduction. In such 
reactions as 

HCl+NaOH->NaCl+H 2 0, 
AgN0 3 +HCl^AgCl+HN0 3 , (169) 

and 

FeCl 3 +3NaOH->Fe(OH) 3 +3NaCl, (173) 

neither oxidation nor reduction occurs. These are called double- 
decomposition reactions. In such reactions no element changes 
its valence. 

If oxidation and reduction take place, the reaction is of a 
distinctly different kind. In such reactions two or more elements 
change valence. 

338. Intensity of Activity of Oxidizing and Reducing Agents. 
— Both oxidizing and reducing agents differ greatly in their 
intensity of activity. For example, manganese dioxide will 
oxidize cold, dilute hydrochloric acid, but oxygen gas will not. 
In consequence we say that manganese dioxide is a stronger 
or more powerful oxidizing agent than oxygen itself. When we 
find, as we may readily do by experiment, that dilute nitric acid 
will oxidize a ferrous salt to a ferric salt and that dilute sulfuric 
acid will not do so, we conclude that nitric acid is a stronger or 
better oxidizing agent than sulfuric acid. When we know that 
hydriodic acid will reduce sulfuric acid, but that hydrochloric 
acid will not do so, we conclude that hydriodic acid is a better 
reducing agent than hydrochloric acid. 

339. Hydrogen Sulfide, H 2 S. — Iron and sulfur unite directly 
at a red heat to form ferrous sulfide, FeS : 

Fe+S->FeS. 

This is a black solid, which is insoluble in water. It reacts 
readily with hydrochloric acid to give ferrous chloride and 
hydrogen sulfide: 

FeS+ 2 HCl->FeCl 2 +H 2 S. 



Oxidation and Reduction 203 

Hydrogen sulfide is a colorless gas, of which water dissolves 

three to four times its own volume. It has a very disagreeable 

odor, resembling rotten eggs, and is extremely poisonous. 

Fatal accidents have often occurred from breathing the gas. 

Hydrogen sulfide is a very powerful reducing agent. In water 

solution it is easily oxidized by atmospheric oxygen, giving sulfur 

and water: 

2H 2 S+0 2 ->2S+ 2 H 2 0. 

A water solution of hydrogen sulfide reacts rapidly with iodine 
to form hydriodic acid (265) and sulfur: 

H 2 S+I 2 ->2HI+S. 

This reaction furnishes a very good practical method for making 
hydriodic acid. We have only to pass hydrogen sulfide gas into 
water containing powdered iodine. When all the iodine has 
been reduced, the solid sulfur can be filtered out, giving a 
clear, colorless filtrate which contains only hydriodic acid and 
water. 

Hydrogen sulfide readily reduces dilute sulfuric acid, which 
is but a very mild oxidizing agent capable of oxidizing only the 
most active reducing agents; the products are sulfurous acid, 
H 2 S0 3 , sulfur, and water: 

H 2 S+H 2 S0 4 ->H 2 S0 3 +S+H 2 0. 

Hydrogen sulfide can be made by the action of sulfuric acid 
on ferrous sulfide, thus: 

FeS+H 2 S0 4 ^FeS0 4 +H 2 S; 

but this is not advisable in practice because of the interaction of 
the sulfuric acid with the hydrogen sulfide. Hydrochloric acid 
is the best acid to use in making hydrogen sulfide. 

Hydrogen sulfide is oxidized by all except the very mildest 
oxidizing agents. As a final example of its behavior, its action 
on ferric salts may be given. These are reduced to ferrous salts, 

thus: 

2FeCl 3 +H 2 S->2FeCl 2 +S+2HCl. 



204 Introduction to General Chemistry 

340. Sulfurous Acid, H 2 S0 3 . — When sulfur burns, it forms 
sulfur dioxide, S0 2 , a colorless gas, with a strong odor: 

S+0 2 ->S0 2 . (286) 

Sulfur dioxide is very soluble in water, with which it unites 
partially to form sulfurous acid: 

S0 2 +H 2 CteH 2 S0 3 . , (286) 

This reaction is reversible; by boiling a solution of sulfurous 
acid the latter can be completely decomposed and all sulfur 
dioxide driven off. 

Sulfurous acid is a reducing agent which, when oxidized, is 
converted into sulfuric acid. It reacts slowly with atmospheric 
oxygen, thus: 

2H 2 S0 3 +0 2 ->2H 2 S0 4 . 

It is rapidly oxidized by chlorine, which is thereby reduced 
to hydrochloric acid: 

H 2 S0 3 +C1 2 +H 2 0->H 2 S0 4 +2HCL 

Manganese dioxide and sulfurous acid react as follows : 
Mn0 2 +H 2 S0 3 ^MnS0 4 +H 2 0. 

Ferric salts are reduced to ferrous salts by sulfurous acid, as 
illustrated by the following equation: 

Fe 2 (S0 4 ) 3 +H 2 S0 3 +H 2 0->2FeS0 4 +2H 2 S0 4 . 

341. Hydrochloric, Hydrobromic, and Hydriodic Acids as 
Reducing Agents. — These acids in water solution can all be 
oxidized, and are therefore to be considered as reducing agents. 
Hydriodic acid is the most easily oxidized of the three, and is 
therefore the best or most powerful reducing agent. It is 
oxidized by atmospheric oxygen (265), which has no action 
whatever on a water solution of hydrochloric acid. The latter 
substance acts as a reducing agent only with respect to the most 
powerful oxidizing agents, such as manganese (234) and lead 
dioxides (235) and potassium permanganate (236). Hydro- 



Oxidation and Reduction 205 

bromic acid is a better reducing agent than hydrochloric acid, 
but not as powerful as hydriodic acid. 

Hydrobromic acid, as a reducing agent, reacts with concen- 
trated sulfuric acid, as an oxidizing agent, as follows, 

2 HBr+H 2 S0 4 ^Br 2 +S0 2 +2H 2 0, 

forming free bromine, sulfur dioxide, and water. 

Hydriodic acid reacts even more vigorously with concentrated 
sulfuric acid. In this case the products vary according to the 
proportions taken, but the reduction of the sulfuric acid may go 
as far as the formation of free sulfur and hydrogen sulfide. The 
possible reactions are represented in the following equations: 

H 2 S0 4 + 2ffl->I 2 + S0 2 + 2 H 2 0, 

H 2 S0 4 +6HI-> 3 I 2 +S+4H 2 0, 

H 2 S0 4 +8HI-> 4 I 2 +H 2 S+ 4 H 2 0. 

It will now be understood why roundabout methods are used 
to prepare hydrogen bromide and hydrogen iodide instead of 
the simple reaction with concentrated sulfuric acid and a salt, 
as is done in the preparation of hydrogen chloride. 

342. Manganese and Its Compounds. — Manganese, Mn, is a 
metallic element which in the free form resembles iron rather 
closely. Its principal ore is the dioxide, Mn0 2 , called by 
mineralogists pyrolusite. Manganese forms a series of salts 
corresponding to the salts of ferrous iron; among such we have 
manganous chloride, MnCl 2 , manganous nitrate, Mn(N0 3 ) 2 , and 
manganous sulfate, MnS0 4 . These salts are pale pink in color 
and are easily soluble in water. The dilute solutions, which are 
almost colorless, give with sodium hydroxide white precipitates 
of manganous hydroxide: 

MnCl 2 + 2NaOH->Mn(OH) 2 + 2 NaCl. 

This hydroxide corresponds to an oxide, MnO: 
Mn(OH) 2 ^MnO+H 2 0. 

In all these compounds except the dioxide, Mn0 2 , manganese is 
bivalent; in the dioxide it is quadrivalent. 



206 Introduction to General Chemistry 

Manganese forms a variety of compounds of a very different 
character from the ones just mentioned; of these the most 
important is potassium permanganate. 

343. Potassium Permanganate, KMn0 4 .— This substance is 
the potassium salt of permanganic acid, HMn0 4 , in which 
manganese acts as an acid-forming element. The salt is made 
from manganese dioxide and potassium hydroxide by compli- 
cated reactions which need not be considered at present. It 
forms dark-purple crystals which dissolve in water to form a 
purple solution having nearly the color of the vapor of iodine. 
It is a very important substance and is one of the most powerful 
of all oxidizing agents. 

We have already learned that potassium permanganate 
oxidizes hydrochloric acid, thus : 

2KMn0 4 +i6HCl->2KCl+ 2 MnCl 2 +5Cl 2 +8H 2 0. (235) 

It can also oxidize almost any substance which is capable of 
being oxidized in solution. Two additional examples may be 
given as illustrations: 

2KMn0 4 + 5H 2 S0 3 H>K 2 S0 4 + 2 MnS0 4 + 2H 2 S0 4 + 3 H 2 0, 

2 KMn0 4 + ioFeS0 4 + 8H 2 S0 4 ->K 2 S0 4 + 2 MnS0 4 + sFe 2 (S0 4 ) 3 + 8H 2 0. 

In the last reaction the sulfuric acid acts neither as a reducing 
nor an oxidizing agent, but is used to keep the solution acid. The 
sulfate radical is not decomposed in the reaction. 

From the foregoing equations it is apparent that when two 
molecules of permanganate change to manganese sulfate or 
chloride, they change the valence of ten atoms (chlorine in 
hydrochloric acid or iron in a ferrous salt) by one unit of valence 
each, or of five atoms by two units of yalence each (sulfur in 
sulfurous acid). This relationship exists because the valence 
of manganese, which is seven in permanganate, changes to two 
in manganese sulfate or chloride: 

344. Chromium and Its Compounds. — The element chro- 
mium, Cr, is a hard metal, resembling iron in appearance. It 
forms a series of salts of which chromic chloride, CrCl 3 , and 
chromic sulfate, Cr 2 (S0 4 ) 3 , are typical examples. Solutions of 



Oxidation and Reduction 207 

chromic salts are either green or violet in color, according to the 
method of preparation. These solutions give with ammonium 
hydroxide bluish precipitates of chromic hydroxide : 

CrCl 3 +3NH 4 OH->Cr(OH) 3 +3NH 4 CL 

The hydroxide when strongly heated gives chromic oxide: 

2Cr(OH) 3 ->Cr 2 3 +3H 2 0. 

It will be seen that chromic salts are analogous to ferric salts 
and that in these compounds chromium is trivalent. 

345. Chromates and Dichromates. — When chromic oxide is 
fused with sodium nitrate or sodium peroxide, sodium chromate, 
Na 2 Cr0 4 , is formed. This is a bright-yellow crystalline salt, 
readily soluble in water. It may be considered as derived from 
chromic acid, H 2 Cr0 4 . Potassium chromate, K 2 Cr0 4 , is also a 
yellow crystalline salt which is readily made by methods similar 
to those that give the sodium salt. 

A solution of potassium chromate, which is bright yellow in 
color, turns deep orange when mixed with sulfuric acid. The 
solution contains potassium dichromate, K 2 Cr 2 7 , which has 
been formed thus : 

2K 2 Cr0 4 +H 2 S0 4 ->K 2 Cr 2 7 +K 2 S0 4 +H 2 0. 

Potassium dichromate forms orange-colored crystals, which 
dissolve in water to form an orange-colored solution. In the 
foregoing reaction we might have expected to get potassium 
hydrogen chromate, KHCr0 4 ; but if this salt is first formed it 
decomposes at once, as follows: 

2KHCr0 4 ->K 2 Cr 2 7 +H 2 0. 

Sodium dichromate, Na 2 Cr 2 7 , orange-colored crystals, can 
be made in a similar manner from sodium chromate. 

346. Chromates and Dichromates as Oxidizing Agents. — 

Solutions of either chromates or dichromates are oxidizing 
agents. More commonly a strongly acid solution is used. Thus 
with hydrogen sulfide an acid solution of potassium chromate 



208 Introduction to General Chemistry 

(potassium dichromate) reacts to form sulfur and chromium 

salts : 

2K 2 Cr0 4 +3H 2 S+ioHCl->4KCl+2CrCl 3 +8H 2 0+3S; 

or if the equation is written for the dichromate we have 
K 2 Cr 2 7 +3H 2 S+8HCl->2KCl+2CrCl 3 +7H 2 0+ 3 S. 

With sulfurous acid as a reducing agent the reaction yields 
sulfuric acid and chromium sulfate: 

2K 2 Cr0 4 +3H 2 S0 3 +2H 2 S0 4 ->2K 2 S0 4 +Cr 2 (S0 4 ) 3 +5H 2 0. 

Apparently two molecules of potassium chromate (or one of the 
dichromate) can cause six units of valence change on other 
atoms, three of sulfur in H 2 S if free sulfur is the product, or 
three of sulfurous acid to form sulfuric acid. 

It is plain that permanganates are more powerful oxidizers 
than chromates or dichromates, since the first can oxidize 
hydrochloric acid and the second cannot, except in very con- 
centrated acid solution. 

347. Hydrogen Peroxide as an Oxidizing Agent. — We have 

already learned (318) that hydrogen peroxide easily decomposes 

into water and oxygen, and that for this reason it acts as an 

oxidizing agent. Its action on hydriodic acid was shown to take 

place thus: 

H 2 2 + 2ffl^I 2 + 2 H 2 0. (320) 

Sulfurous acid is readily oxidized to sulfuric acid : 

H 2 2 +H 2 S0 3 ->H 2 S0 4 +H 2 0. 

Ferrous salts are oxidized to ferric salts as the following 
equation will illustrate : 

H 2 2 +2FeS0 4 +H 2 S0 4 ->Fe 2 (S0 4 ) 3 +2H 2 0. 

Lead forms with sulfur lead sulfide, PbS, a black substance, 
almost insoluble in water. It is obtained as a black precipitate 
by the action of hydrogen sulfide on a solution of a lead salt: 

Pb(N0 3 ) 2 +H 2 S->PbS+2HN0 3 . 



Oxidation and Reduction 209 

Hydrogen peroxide oxidizes lead sulfide to lead sulfate (167) : 
PbS+ 4 H 2 2 ->PbS0 4 +4H 2 0. 

Since lead sulfate is white, the effect of the action is easily seen. 
The blackening of old oil paintings is due to the gradual conver- 
sion of the lead compounds that have served as ingredients of 
the paint into lead sulfide by the action of sulfur compounds 
occurring in the air. Blackened paintings are often restored 
to their original colors by treating them with hydrogen peroxide, 
which converts the black lead sulfide into white lead sulfate. 

It has already been mentioned that animal and vegetable 
substances are bleached by hydrogen peroxide. The exact 
nature of the changes that occur in such reactions is not in general 
known, but it is safe to conclude that they are processes of oxida- 
tion which convert colored into colorless substances. 

348. The Reducing Action of Hydrogen Peroxide. — The 
action of hydrogen peroxide on silver oxide yields free silver, 
and we may say that the silver oxide has been reduced. 

H 2 2 + Ag 2 0-> 2Ag+H 2 0+0 2 . 

Another important reaction of this class is found in the action 
of hydrogen peroxide on potassium permanganate in acid solu- 
tion, which takes place thus: 

5H 2 2 + 2 KMn0 4 -f 3H 2 S0 4 ^ K 2 S0 4 + 2 MnS0 4 +8H 2 0+ 50 2 . 

The products are the colorless solution of the sulfates of potas- 
sium and manganese in addition to free oxygen. 

349. Hypochlorous Acid, HCIO. — It is probable that chlorine 
reacts reversibly with water in which it is dissolved to form 
hydrochloric acid and hypochlorous acid, HCIO, thus: 

Cl 2 +H 2 O^HCl+HC10. 

Since this is a reversible reaction, all four substances are con- 
tained in equilibrium in a solution of chlorine in water. Hypo- 
chlorous acid is very unstable, that is, it easily decomposes, 
and for this reason it cannot be obtained except in the form of a 
dilute water solution. It has only very weak acid properties 



210 Introduction to General Chemistry 

and cannot even decompose calcium carbonate, which is acted 
upon by almost all other acids. As is well known, hydrochloric 
acid reacts with calcium carbonate as follows : 

2HCl+CaC0 3 ->CaCl 2 +C0 2 +H 2 0. (163) 

As a matter of fact, when calcium carbonate is added to chlorine 
water it reacts as follows : 

CaC0 3 + 2 Cl 2 +H 2 0-> 2HCIO+ CaCl 2 + C0 2 . 

From the resulting solution hypochlorous acid mixed with much 
water vapor can be driven off by cautious heating; ' the condensed 
vapor forms a dilute solution of hypochlorous acid. This reac- 
tion seems to prove that chlorine and water react to form hydro- 
chloric and hypochlorous acids. 

350. Hypochlorites. — If chlorine is passed into a cold, dilute 
solution of sodium hydroxide, sodium chloride and sodium hypo- 
chlorite, NaCIO, are formed: 

Cl 2 + 2NaOH ->NaCl+NaC10+H 2 0. 

This is exactly what we should expect if both acids which result 

from the action of chlorine on water are neutralized by the 

sodium hydroxide. Chlorine and potassium hydroxide react 

similarly : 

Cl 2 +2KOH->KCH-KC10+H 2 0. 

351. Bleaching Powder. — The action of chlorine gas on solid 
slaked lime, calcium hydroxide, takes place thus: 

2 Cl 2 +2Ca(OH) 2 ->CaCl 2 +Ca(C10) 2 +2H 2 0. 

The product of the reaction is a white powder known as chloride 
of lime or bleaching powder. It is a mixture of calcium chloride 
and calcium hypochlorite. It is extensively used in the bleaching 
of cotton goods and for a variety of other purposes. Before 
taking up the chemical behavior of hypochlorous acid and hypo- 
chlorites it will be of interest to consider the formation of these 
substances from the standpoint of oxidation and reduction. 



Oxidation and Reduction 



211 



352. The Oxidation Products of Chlorine. — By the action of 
chlorine gas on dry mercuric oxide, HgO, chlorine monoxide, 
C1 2 0, a brownish-yellow gas, is obtained. It is obvious that in 
this reaction the chlorine has been oxidized. Now this oxide 
of chlorine unites with water to form hypochlorous acid, 

C1 2 0+H 2 0->2HC10. 

The relation between chlorine monoxide and hypochlorous 
acid is similar to that between sulfur dioxide and sulfurous acid: 

S0 2 +H 2 0->H 2 S0 3 . 

Chlorine and sulfur also show similar behavior in that each forms 
compounds with hydrogen and metals, namely chlorides and 
sulfides. 

353. The Formation of Chlorates. — Hypochlorites are very 
unstable salts. A warm, concentrated solution of sodium hypo- 
chlorate changes more or less rapidly into sodium chloride and 
sodium chlorate, NaC10 3 , according to the equation, 

3 NaC10 -> 2 NaCl+NaC10 3 . 

Potassium hypochlorite changes in a similar fashion, yielding 
potassium chlorate, KC10 3 . 

Sodium and potassium chlorates are powerful oxidizing 
agents, since they contain large proportions of easily liberated 
oxygen. When the dry crystals are heated they decompose 
finally into chlorides and oxygen: 

2KC10 3 ->2KCl+ 3 2 . 

This reaction takes place in two stages (306) . The first change 
gives rise to a perchlorate, KC10 4 , thus: 

ioKC10 3 ->4KCl+6KC10 4 +30 2 . 

Until recently potassium chlorate was used extensively, and 
sodium chlorate was rarely seen. The reason was twofold: in 
the first place potassium chlorate was made very largely in 
Germany, where potassium compounds are cheap on account 
of the immense potash deposits found in that country; and in 



212 Introduction to General Chemistry 

the second place sodium chlorate, being more soluble, is more 
difficultly purified than the potassium salt. Since the war began 
there has been a shortage of potash, because no other country 
besides Germany has much easily accessible potash. As a 
consequence the manufacture of sodium salts has been stimu- 
lated, and since 191 5 there has been an abundant supply of 
sodium chlorate. This can be used advantageously in place of 
potassium chlorate for nearly all purposes. 

354. Chloric Acid and Chlorine Dioxide. — Potassium and 
sodium chlorates are salts of chloric acid, HC10 3 . This is a very 
unstable acid, which is known only in dilute solution. Upon 
evaporation of the solution the acid decomposes, giving chlorine 
dioxide, C10 2 , and other products. 

If a few drops of concentrated sulphuric acid are poured on 
a small crystal of sodium chlorate in a dry test tube, a yellow 
gas forms, which explodes with violence a few seconds later. 
This dangerous experiment should be performed with great 
caution. The yellow gas is chlorine dioxide, C10 2 , which was 
formed by the decomposition of the chloric acid set free, thus: 

NaC10 3 +H 2 S0 4 ->NaHS0 4 +HC10 3 . 

The explosion of chlorine dioxide is due to decomposition into 
its elements: 

2C10 2 ->Cl 2 +20 2 . 

Chloric acid is a powerful oxidizing agent. For example, it 
changes lead sulfide to lead sulfate (167). This operation is 
usually carried out by adding a few crystals of sodium chlorate 
and dilute hydrochloric acid to the black lead sulfide. The 
dark color is seen to change slowly to the white of the sulfate: 

3 PbS+ 4 HC10 3 -> 3 PbS0 4 +4HCL 

355. Perchlorates and Perchloric Acid. — Perchlorates are 
formed by heating chlorates gently (306, 353). 

Sodium perchlorate, NaC10 4 , and potassium perchlorate, 
KC10 4 , are white crystalline salts. They decompose completely 
into chlorides and oxygen at dull-red heat. For example, 

NaC10 4 ->NaCl+20 2 . 



Oxidation and Reduction 213 

Ammonium perchlorate, NH 4 C10 4 , is made by neutralizing 
perchloric acid, HC0 4 , with ammonia. It is used as an oxidizing 
agent and as a very powerful explosive. 

When powdered sodium or potassium perchlorate is mixed 
with concentrated sulfuric acid and cautiously heated in a small 
retort (104, Fig. 24), perchloric acid, HC10 4 , is distilled from the 
mixture. This experiment should not be made by the student, 
as it might result in an explosion in unskilled hands. 

NaC10 4 +H 2 S0 4 -^NaHS0 4 +HC10 4 . 

Perchloric acid is a colorless liquid. It is a violent oxidizing 
agent, as shown by the fact that a drop of the acid will set fire to 
filter paper. The diluted acid is now coming into use in labo- 
ratories as an oxidizing agent, and also for the purpose of pre- 
cipitating potassium perchlorate in the quantitative analysis 
of potassium. 



CHAPTER XVI 



HEAT AND ENERGY 

356. Heat of Combustion. — Since coal, wood, and fuel gas 
are burned ordinarily in order to produce heat rather than as 
a means of obtaining their products of combustion, carbon 
dioxide and water, it becomes a matter of importance to discover 

how much heat is produced 
in the burning of a known 
weight of a given substance. 
The unit of heat is the 
calorie (in), which is the 
amount of heat required to 
raise the temperature of one 
gram of water one degree centi- 
grade. The amount of heat 
produced by the burning of 
one formula weight of a pure 
substance is called its heat 
of combustion. The heat of 
combustion of a solid is de- 
termined by burning a known 
weight of it within an appa- 
ratus of special design, called 
a bomb calorimeter. 

357. The Bomb Calorimeter. — This apparatus, illustrated in 
Fig. 44, consists of a heavy-walled metallic bomb with a gas- 
tight cover, surrounded by a vessel of water. The latter is con- 
tained in a larger vessel with walls of heat-insulating material. 
A weighed amount of substance whose heat of combustion is 
to be found is placed in the crucible of the bomb, which is filled 
with oxygen gas. The substance is then ignited by heat from 
a wire which carries an electric current. The temperature of 
the water surrounding the bomb is measured accurately before 




Fig. 44 



214 



Heat and Energy 



215 



and after the burning, and the number of calories of heat pro- 
duced is calculated from the rise of temperature and the weight 
of water actually heated, plus the water equivalent of the bomb, 
etc. The water equivalent is the amount of water which has 
the same heat capacity as the bomb and other heated parts of 
the apparatus. Some typical results of measurements of heats 
of combustion are shown in Table XI. The values are given 
to the nearest hundred, since this is about the limit of accuracy 
in such measurements. 

TABLE XI 



Substance 


Calories per gram 


Formula 


Heat of Combustion 


Carbon 

Hydrogen 


8,130 
34,400 

2,200 
11,900 

2,430 


C=I2g. 
H 2 = 2g. 

S =32 g. 

C 2 H 2 = 26g. 

CO =28g. 


97,600 
68,800 
70,400 
315,400 
68,200 


Sulfur 


Acetylene 

Carbon monoxide 







Since one formula weight of a gaseous substance has a volume 
22.4 liters, the heats of combustion of H 2 , C 2 H 2 , and CO are 
the amounts of heat produced in the burning of equal volumes 
of these gases. It will be seen that the heat of combustion of 
C 2 H 2 is very large (nearly five times that of hydrogen). This 
accounts in part for the very high temperature of the oxyacety- 
lene flame (315). 

358. The British Thermal Unit, B.T.U.— In engineering 
practice quantities of heat are measured in British Thermal 
Units (B.T.U.) instead of in calories. This unit is the amount 
of heat required to raise the temperature of one pound of water one 
degree Fahrenheit. Since one pound equals 453 g., and i° F. = 5/9 
of i° C, it follows that 1 B.T.U.= 252 calories. The heat pro- 
duced in burning coal, coke, and fuel gas is called its calorific 
power. It is usually stated in terms of B.T.U. per pound of 
fuel. 

359. Composition and Calorific Power of Fuel. — Since the 
value of fuel is directly dependent on its calorific power, the 
testing of fuel is a matter of great practical importance. In 
testing coal it is customary to determine the moisture, volatile 



2l6 



Introduction to General Chemistry 



matter, "fixed carbon," and ash in addition to the calorific 
power. The "fixed carbon" is the non- volatile residue left 
when all volatile matter is driven off at a bright-red heat in the 
absence of air, less the ash contained therein. The calorific 
power is usually expressed in terms of B.T.U. per pound of fuel, 
or per cubic foot in the case of gases. Table XII gives some 
results for a variety of solid fuels. 

TABLE XII 





Percentage Composition 


Calorific Power 


Kind of Fuel 


Volatile 
Matter 


Fixed 
Carbon 


Ash 


Calories 
per gram 


B.T.U. 

per pound 


Lackawanna anthracite coal 
Pocahontas coal 


5 
18 

35 
o-5 
38 


84 
74 
50 
90 

5i 


11 

7 
6 

9 
4 

0.4 
0.4 


7,724 
8,760 
8,080 
7,900 
7,200 
4,600 
5,000 
11,520 


13,900 
15,680 
14,540 
14,200 


Indiana bituminous coal . . . 
Coke 


Lignite 


13,000 
8,300 
9,IOO 

20,736 


Oak wood 


Pine wood (resinous) 






Crude petroleum 















Table XIII gives the calorific power of some typical fuel 

gases. 

TABLE XIII 

Calorific Power in B.T.U. per Cubic Foot 

Kokomo, Indiana, natural gas 1 ,000 

Pittsburgh, Pennsylvania, natural gas 1,150 

Coal gas 650 

City of Chicago gas , 600 

360. The Evaporation of Water and the Production of 
Steam. — We can easily calculate the amount of fuel theoretically 
needed to change water at ordinary temperature into steam. 
If one gram of water at 20 is heated to ioo°, 80 calories of heat 
are required, and in addition 540 calories are needed to change 
this hot water into steam. The total is 620 calories. Since 
the burning of one gram of coal produces about 8,000 calories, 
if all this heat were utilized it would be sufficient to evaporate 
(8,ooo-f- 620) 13 g. of water. In practice much heat is lost to 
the surroundings, as well as in the hot smoke which goes up the 



Heat and Energy 217 

smokestack. Engineers consider that it is good practice to 
evaporate 8 g. of water with 1 g. of coal. Therefore one pound 
of good coal will change 8 lb., or about 1 gal. (8.3 lb) of water 
at ordinary temperature into steam at ioo°. 

361. Heat of Reaction and Heat of Formation. — We have 
already frequently observed that numerous reactions other than 
combustions in oxygen (air) produce much heat. Among such 
are the reactions of chlorine with hydrogen (244), phosphorus 
(247), antimony (246), and turpentine (248); and water with 
sulfuric acid (93), potassium (106), and calcium oxide (150). 
The heat produced in these and other reactions may be measured 
in suitably constructed calorimeters and the results expressed 
most conveniently by stating the amount of heat given out in 
the reaction of formula weights of the uniting substances; or 
in the formation of one formula weight of the product. Thus 
the heat of reaction of CaO and H 2 may be written 

CaO+H 2 0->Ca(OH) 2 +5,ioo cal. 
and the heat of formation of water from its elements 
H 2 +|0 2 ->H 2 0+68,8oo cal. 

362. Heat of Neutralization. — The union of acids and bases 
to form salts and water always gives out heat. In fairly dilute 
solutions the amount of heat given out when one formula weight 
of water is so formed is almost exactly the same for many acids 
and bases. For example. 

HC1, NaOH = 13,700 cal. 
HC1, KOH =13,700 " 
HNO3, NaOH = 13,700 " 
HNO3, KOH =13,700 " 

This regularity is indeed striking and must mean close similarity 
in the processes of these reactions. How chemists interpret this 
phenomenon will be considered in chapter xviii. 

363. The Law of Constant Heat Summation. — Let us now 
consider the following question: If equal quantities of a given 
substance can be changed into the same product by two different 



218 Introduction to General Chemistry 

ways, will the amounts of heat produced be the same in the two 
cases? Carbon, for example, gives carbon dioxide when it is 
burned, 

C+0 2 ->C0 2 , 

but in a deficiency of oxygen the product is carbon monoxide, 

2C+0 2 ->2CO. 

Carbon monoxide is a colorless gas which burns readily, giving 
carbon dioxide, 

2CO+0 2 ^2C0 2 . 

Therefore it is possible to change given weights of carbon and 
oxygen into carbon dioxide in two different ways. The heats 
of combustion are as follows.: 

First Way 
C+§0 2 ->C0 +29,400 cal. 
CO+i0 2 ->C0 2 +68,2oo cal. 



Sum 97,600 cal. 

Second Way 
C+0 2 ^>C0 2 +97,6oo cal. 

These results show that if 1 2 g. of carbon (C = 1 2) unite with 
$2 g. of oxygen (0 2 = 32 and |0 2 = 16) the total heat produced is 
the same no matter in which way the union occurs. 

Another illustration is found in the formation of a solution 
of ammonium chloride, NH 4 C1, from NH 3 and HC1 gases. 
This reaction can take place in two ways : 

First Way 

NH 3 (gas)+HCl (gas)->NH 4 Cl (solid) +42,000 cal. 
Heat absorbed in dissolving the NH 4 C1 in water = — 3,900 cal. 

Excess of heat produced over heat absorbed = 38,100 cal. 

Second Way 

Heat of solution of NH 3 in water = 8,400 cal. 

Heat of solution of HC1 in water = 17,300 cal. 

Heat of neutralization of the two solutions = 12,400 cal. 



Total heat produced = 38,100 cal. 



Heat and Energy 219 

Innumerable cases like the two here given in illustration have 
led to the Law of Constant Heat Summation (Law of Hess). 
The heat produced or absorbed in the change of given substances 
into the same final products {in the same physical state) is the same, 
by whatever way the changes occur. 

That the heat of a given reaction is dependent on the physical 
state of the reacting substances and products is illustrated by 
the following example: 

CaO+H 2 (liquid)->Ca(OH) 2 (solid) + 15,100 cal. 
Ca0+H 2 (ice) ->Ca(0H) 2 (solid) + 13, 7 00 cal. 



Difference =1,400 cal. 

The difference, 1,400 cal., is due to the fact that it requires this 
amount of heat to change one formula weight of ice into water 
(18X79 = 1,422) (118). 

364. Heat Produced in Slow Oxidation. Spontaneous Com- 
bustion. — Numerous experiments have proved that the amount 
of heat formed in a given reaction is just the same whether the 
change takes place slowly or rapidly. The decay of wood leads 
ultimately to the production of carbon dioxide and water, the 
same products as those formed when wood is burned. During 
the decay of wood, heat is produced so slowly that its formation 
is usually not perceptible by ordinary observation. Coal also, 
when exposed to the air, slowly oxidizes. In so doing it often 
loses an appreciable part of its heating value before it is burned. 
The depreciation on this account in the value of stored coal is a 
matter of considerable importance. 

If coal (especially bituminous coal) in small lumps and con- 
taining much dust is heaped in immense piles, such as are seen 
in coal yards, the heat produced by the slow oxidation does not 
escape readily from the bottom layers of the pile. The result 
is a gradual rise of temperature. At the higher temperature 
oxidation and therefore heat production go on still faster, since 
usually enough air can diffuse in to keep up the supply of oxygen. 
Finally the temperature may rise so high that the pile of coal 
actually takes fire at the surface, where there is of course an 
unlimited supply of oxygen. Fire originating in this way is 



220 Introduction to General Chemistry 

said to be due to spontaneous combustion. The loss of coal 
through such fires was a very serious feature of the "coal famine" 
of 191 7-18. Some smoke is seen issuing from the majority of 
large piles of low-grade coal in the Chicago district, thus indi- 
cating more or less fire beneath. It is almost impossible to 
extinguish fire in a very large coal pile. The best way to prevent 
serious rise of temperature in coal piles is to provide numerous 
air shafts in the pile, by means of which warm air can escape. 
This does not entirely prevent oxidation but keeps the tempera- 
ture down to a point where the oxidation is not dangerously 
fast. 

It is a popularly known fact that " greasy" rags will often 
catch fire spontaneously. As a matter of fact such fires originate 
usually in rags soaked in oils used in paint or varnish, especially 
linseed oil or turpentine. The "drying" of paint and varnish 
is not a process of evaporation as much as one of oxidation of 
the oil used. These paint and varnish oils readily unite with 
oxygen to form solid products. In this process heat is produced. 
In a pile of rags, etc., covered with such oils sufficient rise of 
temperature may occur to cause spontaneous combustion. For 
this reason greasy rags, etc., should never be left where they 
can do damage if they take fire. 

365. Dust Explosions. — When the air is filled with the dust of 
coal, wood, flour, or other combustible substance a flame will 
often start a combustion which will spread with explosive 
rapidity. Appalling explosions have occurred from such causes 
in coal mines, wood-working factories, and flour mills. Even 
dust which is at rest in such places is blown into the air by the 
on-coming explosion wave and is thus changed to an explosive 
dust and air mixture. It is easy to see that a dust explosion is 
due to the extremely rapid burning of minute particles, each 
surrounded by an abundance of oxygen. Dust explosions are 
best prevented by keeping mines, mills, etc., free from accumula- 
tions of dust. 

366. Modes of Heat Production in Physical and Chemical 
Changes. — We have now learned that heat is produced (or 
absorbed) in a variety of physical and in all chemical changes. 



Heat and Energy 221 

The following seven modes of heat production (or absorption) 
have been studied: 

1. Latent heat of fusion (melting) (118). 

2. Latent heat of evaporation (115). 

3. Heat of solution (127). 

4. Heat of combustion (356). 

5. Heat of formation (361). 

6. Heat of reaction (361). 

7. Heat of neutralization (362). 

The first three modes have to do with physical changes of the 
sort known as changes of state ; the last four are due to chemical 
changes. All changes of state and many chemical changes are 
reversible processes. In every reversible process, if heat is 
given out when the -change proceeds in one direction, heat is 
absorbed in equal amount when the change proceeds to an equal 
extent in the opposite direction. A change which results in the 
production of heat is called an exothermic change ; one in which 
heat is absorbed is an endothermic change. 

367. Heat Production and Equilibrium. — In chapter xiii 
(288) the effect of temperature on equilibrium was discussed 
briefly. With respect to the change of solubility it was stated 
that raising the temperature causes that change of solubility to 
occur which involves an absorption of heat. We also saw (288) 
that for chemical equilibrium raising the temperature causes 
the equilibrium to shift in the direction that involves an absorption 
of heat. These laws are entirely general and apply to all 
changes of state and all chemical changes. 

In the shift of equilibrium which occurs with change of 
temperature the fraction of the reacting substances transformed 
to new products is determined, in a given case, by the change 
of temperature (measured in degrees). The amount of heat 
(in calories) absorbed (if the temperature is raised) or given out 
(if the temperature is lowered) is determined by the amount of 
material transformed. An example will make the matter clearer. 

Hydrogen and iodine vapor react partially in the neighbor- 
hood of 400 to give hydrogen iodide (264, 281, 288): 
H 2 +I 2 ^2HI+i,ooo cal. 



222 Introduction to General Chemistry 

This equation means that the formation of two formula weights 
of HI from H 2 and I 2 (vapor) at about 400 takes place with 
the liberation of 1,000 cal. of heat or 500 cal. for each formula 
weight of HI produced. The following table shows the propor- 
tions of molecules in the equilibrium mixture at 370 and 440 : 





H, 


I, 


HI 


Total 


370 


IOO 


IOO 


800 


1,000 


440 


I IO 


no 


780 


1,000 



We see that if the temperature is raised from 370 to 440 , 20 
molecules of HI out of a total of 1,000 molecules (2 per cent of 
the whole) change into H 2 and I 2 . If the total amount of 
material in the mixture is that resulting from one formula weight 
each of H 2 and I 2 (equivalent to two formula weights of HI), 
and if 2 per cent of the whole number of molecules change into 
H 2 and I 2 , the heat absorbed is 0.02 X 1,000 cal. = 20 cal. 

368. Work and Energy. — The terms work and energy have 
very definite meanings in science. The subject of physics is 
largely concerned with these very important matters; and since 
it is assumed that the student has already studied physics, an 
elementary discussion of these very important topics is unneces- 
sary. We may, however, briefly summarize some of the more 
prominent points. The typical example of work in the physical 
sense is the lifting of a weight. The scientific unit of work is 
the gram centimeter, which is the work required to lift one gram 
one centimeter. The amount of work done in lifting a weight is 
the product of the force required (which in this case is equal 
to the weight in grams) and the vertical distance measured in 
centimeters. Thus the lifting of 600 g. to a height of 30 cm. 
requires the doing of 600X30 = 18,000 g.cm. of work. The 
weight of 600 g., having been lifted 30 cm., could do work to 
the extent of 18,000 g.cm. in descending 30 cm. It is said to 
have the power to do this amount of work. Now power to do 
work is called energy, and therefore it has 18,000 g.cm. of 
energy. Two kinds of energy are recognized: potential energy, 
as possessed by a weight which may do work on descending, and 
kinetic energy, or the energy of a body in motion. It requires 



Heat and Energy 



223 



work to set a body in motion, and conversely a body in motion 
is able to do work. 

369. The Mechanical Equivalent of Heat. — Heat is also a 
form of energy, because heat is able to do work. A steam engine 
is merely a machine which converts the heat of burning coal into 
kinetic energy. The change of kinetic energy into heat may be 
observed on every hand: anything that restrains or stops the 
motion of a moving body converts part or all of its kinetic energy 
into heat. We measure energy in gram centimeters and heat 
in calories, and if heat is a form of energy then the calorie, like 
the gram centimeter, must be an energy unit. It will at once 




w 



Fig. 45 



be asked: Do these units represent equal amounts of energy? 
In other words, will one gram centimeter of work produce one 
calorie of heat? If not, how many gram centimeters are required 
to produce one calorie? This question was first answered by 
Joule in 1840. 

370. Joule's Experiment. — In Joule's experiment, with 
apparatus shown in Fig. 45, a weight, W, attached to a cord 
wound on a cylinder, in slowly descending turns a stirrer which 
is surrounded by water in a calorimeter, C. The water, which 
restrains the motion of the stirrer, becomes warmer, owing to 
the change of work into heat. The amount of work in gram 
centimeters done in heating the water is the product of the mass 
in grams of the weight and the distance of its descent in 



224 Introduction to General Chemistry 

centimeters. The amount in calories of heat produced is the 
product of the rise in temperature in degrees C. and the mass, 
in grams, of water plus the water equivalent of the heated parts 
of the calorimeter. 

By means of this apparatus Joule found pretty closely the 
number of gram centimeters of work equivalent to one calorie 
of heat. More refined work since then has shown that one 
calorie is equal to 42,700 g.cm. This ratio is called the mechanical 
equivalent of heat. This means, for example, that one gram 
falling 42,700 cm. (a little over a quarter of a mile) produces one 
calorie. 

371. The Conservation of Energy. — At the time Joule began 
his experiments in 1840 it was not at all clear that the amount 
of heat produced by a given amount of work (kinetic or potential 
energy) was definite. It seemed possible, if not probable, that 
different modes of changing work into heat would give different 
values for the mechanical equivalent. So Joule used not only 
the method and apparatus already described but also two others. 
His three methods and the mechanical equivalent of one calorie 
were as follows: (1) stirring water in a brass vessel with a brass 
paddle, 42,400 g.cm.; (2) stirring mercury in an iron vessel 
with an iron paddle, 42,500 g.cm.; (3) rubbing two iron rings 
together under mercury, 42,500 g.cm. 

The very close agreement of the results of the three experi- 
ments led Joule to conclude that the amount of heat produced 
by a given amount of work is always the same, by whatever way the 
work is changed into heat. This result has been amply confirmed 
by all later experiments and experience. When work of any 
kind (mechanical energy, either kinetic or potential) is changed 
into heat there is nojreal loss or destruction of energy, since the 
heat produced is also energy in another form and exactly equal 
in amount to the work done in producing it. This conclusion 
is concisely stated in the Law of the Conservation of Energy: 
Energy is indestructible. 

Just as the law of the conservation (indestructibility) of 
matter (21) is the foundation stone of the science of chemistry, 
so, similarly, this law of the conservation (indestructibility) 



Heat and Energy 225 

of energy is the solid rock upon which the whole structure of the 
science of physics rests. 

372. Other Forms of Energy. — We have defined the term 
energy as the power of doing work; and since heat is also a form 
of energy, we might extend the definition so as to read: Energy 
is the power to do work or produce heat. According to this defini- 
tion of energy it is obvious that light and even sound and espe- 
cially electric currents are also forms of energy, since each of 
these by appropriate means can produce work or heat. 

373. Chemical Energy. — For the chemist an important ques- 
tion now arises : What shall be said of the source of energy that 
produces the great heat of a burning substance? This question 
is somewhat like the one, What is the source of energy of a 
" wound-up" watch spring? To wind up the spring a certain 
amount of work must be done. Is it not reasonable to say that 
the energy used in winding up the spring has been " stored up" 
in the coiled spring? If so, we may say that this energy is 
changed into potential energy, just as we say that the energy 
required to lift a weight is changed into potential energy and 
can be regained as useful work then the weight is allowed to 
descend. Reasoning somewhat similarly, we may conclude that 
the energy given out as heat in the burning of hydrogen, for 

which we have 

H 2 +J0 2 ^H 2 0+68,8oo cal., 

comes from some form of potential energy which has been stored 
up in the two gases. This conclusion is rendered highly probable 
by reason of the fact that by means of an electric current (elec- 
trical energy) we can decompose water into hydrogen and oxygen. 
Since the electrical energy disappears and very little heat is 
formed, we may very reasonably conclude that it has been 
changed into some sort of potential energy stored up in the two 
gases formed from the water. The form of potential energy 
stored up in chemical substances and liberated when they react 
is called chemical energy. 

374. The Sun as a Source of Energy. — It will be interesting 
to trace some familiar form of energy through various trans- 
formations back to its source. Take, as an example, the energy 



226 Introduction to General Chemistry 

given out as light and heat by an electric lamp. The energy 
comes to the lamp as an electric current having electrical energy. 
This electrical energy was produced in a dynamo or generator, 
the armature (the moving part) of which was turned by a steam 
engine. The kinetic energy of the engine was derived from hot, 
compressed steam produced from water by the burning of coal 
which has resulted from the slow transformation of vegetable 
matter. 

Plants derive nearly all of their substance from water and 
the carbon dioxide of the air under the influence of the light and 
heat of the sun. A great deal of energy is taken up by plants 
as light and heat and is stored as chemical energy in the sub- 
stances composing them, as well as in the oxygen which is set 
free by the growing plant. Recapitulating, we see that the 
light and heat from the sun are changed by growing plants into 
chemical energy; this energy is largely conserved when plants 
are changed into coal. When the coal burns, its chemical 
energy, supplemented by that of the oxygen of the air, is changed 
into heat, which is in turn changed into kinetic energy in the 
steam engine. The kinetic energy of the engine is then changed 
by a dynamo into electrical energy, and the latter produces in 
the lamp heat and light. 



CHAPTER XVII 
THE IONIC HYPOTHESIS 

375. The Ionic Hypothesis. — This chapter will treat of the 
properties and behavior of acids, bases, and salts and aims- to 
show how a supposition called the ionic hypothesis furnishes a 
satisfactory explanation of many facts. 

376. The Two Parts of a Salt. — It must have been noticed 
that a salt is made up of two parts, the metallic or basic part and 
the non-metallic or acidic part. The latter may be an element 
like chlorine in sodium chloride; or it may be a radical (147) 
like S0 4 , which is contained in every sulfate. The name of a 
salt always indicates the parts of which it may be considered as 
being made up. Thus potassium nitrate, KN0 3 , is composed 
of potassium and nitrate radical, N0 3 ; and calcium carbonate, 
CaC0 3 , of calcium and carbonate radical, C0 3 . 

377. The Two Parts of an Acid. — Every acid may also be 
considered as made up of two parts, one of which is hydrogen 
and the other the characteristic acid radical of that acid. For 
example, sulfuric acid, H 2 S0 4 , may be considered to consist of 
hydrogen and sulfate radical, S0 4 ; and phosphoric acid, H 3 P0 4 , 
to consist of hydrogen and phosphate radical, P0 4 . For this 
reason S0 4 and P0 4 may be called acidic radicals. Dilute 
solutions of pronounced acids have a sour taste. Since hydro- 
gen is the only constituent which all acids have in common, we 
may reasonably conclude that the sour taste is due to the H 
radical. 

378. The Two Parts of a Base. — A base is usually the 
hydroxide of a metallic element, and it may therefore be con- 
sidered as made up of two parts, the metal and the hydroxyl 
radical, OH. Thus sodium hydroxide, NaOH, consists of 
sodium and hydroxyl radical, OH. Ammonium hydroxide, 
NH 4 OH, may be considered as made up of ammonium radical, 
NH 4 , and hydroxyl. Consequently the ammonium radical may 

227 



228 Introduction to General Chemistry 

be called a basic radical. It is the only basic radical that we 
have studied, although many others are known. 

379. The Process of Neutralization. — The two following 
equations represent typical cases of neutralization: 

NaOH+HCl->NaCl+H 2 ; 
NH 4 OH+HN0 3 ->NH 4 N0 3 +H 2 0. 

We notice that in each case the salt which is formed is made up 
of two parts, one of which comes from the base, the other from 
the acid. In each case water, whose formula may be written 
HOH, is also formed. We might call water hydrogen hydroxide 
and think of it as being made up of two parts hydrogen and 
hydroxyl radical. The process of neutralization consists, there- 
fore, merely of an exchange of partners, so to speak, on the part 
of the base and the acid. 

As a matter of fact, not only can neutralization be represented 
in this way, but most reactions in water solution between acids, 
bases, and salts which do not involve oxidation or reduction 
may be regarded as an exchange of the partners of the reagents 
initially used. This will be made clear by the following examples. 

380. Reactions of Barium Salts with Sulfates. — If we add 
dilute sulfuric acid to a dilute solution of barium chloride a white 
precipitate of barium sulfate is formed, 

H 2 S0 4 +BaCl 2 ->BaS0 4 +2HCl. 

A precipitate of barium sulfate also results when a solution 
of any barium salt is added to a solution of any sulfate, as 
illustrated by the following equations: 

K 2 S0 4 +Ba(N0 3 ) 2 ^BaS0 4 +2KN0 3 , 
CuS0 4 +BaBr 2 ->BaS0 4 +CuBr 2 . 

This is so generally true that the formation of a precipitate 
of barium sulfate upon the addition of a solution of a barium 
salt to some other solution shows that this second solution con- 
tains the sulfate radical, S0 4 , in the form either of a sulfate or 
of sulfuric acid. We say therefore that the formation of a 
precipitate of barium sulfate when a solution of a barium salt is 
added to another solution is a test for the sulfate radical. It is 



The Ionic Hypothesis 229 

important to note that it is the S0 4 radical, and not sulfur or 
oxygen or a combination of the two in some other proportion, 
that responds to this test. A solution of hydrogen sulfide, 
H 2 S (339), which may be considered as being made up of two 
parts, hydrogen and sulfur, does not give a precipitate of any 
sort with a solution of a barium salt. Furthermore, pure dilute 
sulfurous acid, H 2 S0 3 (340), which is made up of hydrogen and 
sulfite radical, S0 3 , does not give a precipitate when mixed with 
a barium salt solution. 

381. Reactions of Simple Lead Salts. — Lead sulfate, PbS0 4 , 
is also a white insoluble salt. If we add a solution of any 
simple lead salt to a dilute solution of sulfuric acid or any soluble 
sulfate, we obtain a white precipitate of lead sulfate, 

Na 2 S0 4 +Pb(N0 3 ) 2 ^PbS0 4 +2NaN0 3 . 

In this case, just as in the precipitation of barium sulfate, it is 
the sulfate radical, S0 4 , which has united with the lead to form 
the precipitate. 

It is also of equal interest to note that if the nitrate of barium 
or of lead is used, the nitrate radical, N0 3 , is left in combination 
with the basic element or radical which was originally combined 
with the sulfate radical. 

382. The Reaction of Silver Salts with Chlorides. — We have 
already learned that a solution of silver nitrate reacts with 
hydrochloric acid or a chloride to give a white precipitate of 
silver chloride : 

AgN0 3 +NaCl->AgCl+NaN0 3 . (169) 

A solution of any simple silver salt reacts similarly with hydro- 
chloric acid or any chloride, so that we may think of the reaction 
as characteristic and call it a test for silver salts. This reaction 
is specifically that of the chloride radical; for if we add silver 
nitrate solution to a solution of potassium chlorate, KC10 3 
(353)? no apparent change is observed; certainly no silver 
chloride is formed, otherwise the latter, being insoluble, would 
separate out as a white precipitate. This shows that chlor- 
ine in the chlorate radical, C10 3 , behaves entirely differently 



230 Introduction to General Chemistry 

from chlorine in the form of a chloride. We also find that 
solutions of perchlorates, of which potassium perchlorate, 
KCIO4 (355), is an example, -do not give precipitates with solu- 
tions of silver salts. It is possible to make both silver chlorate, 
AgC10 3 , and silver perchlorate, AgC10 4 , by methods which we 
need not consider at present, and it is found that these salts are 
entirely different from silver chloride, and that both are easily 
soluble in water. 

This brief review of reactions, most of which have already 
been studied, is sufficient to illustrate the subject in hand, but 
many other examples of the same principle will be found in the 
previous chapters. 

383. Summary and Conclusions. — The observations which 
we have made are typical for all acids, bases, and salts. Each 
may be shown to be made up of two parts. In the examples we 
have studied these are, on the one hand, hydrogen, a metal, or 
the ammonium radical, and, on the other, hydroxyl, a halogen, 
sulfur, or an acid radical. Hydrogen is one of the two parts of 
every acid, and hydroxyl one of the two parts of every base. 
In chemical reactions between acids and bases, acids and salts, 
bases and salts, and between two salts (where oxidation and 
reduction do not occur) the two substances simply exchange 
parts. This kind of' chemical change is called double decompo- 
sition (337) or metathesis. The chemical reactions which acids, 
bases, and salts give are in reality only the reactions of their parts. 

Finally it should be noted that the recombination of these 
parts always takes place between the basic or metallic part on 
the one hand and the acidic or the non-metallic part on the 
other. Double decompositions in water solutions never give 
compounds such as KNa or C1S0 4 . This is a striking observa- 
tion, and the fact that we have as yet no explanation for it 
warns us at once that we must go farther in our observations 
to understand even the most commonplace of these reactions. 

384. Double Decomposition and Electrical Conductivity.— 
Along with the ability to undergo double decompositions, acids, 
bases, and salts in their water solutions have the property of 
conducting the electric current. If we set up a battery, a 



The Ionic Hypothesis 



231 



galvanometer, and a salt solution in the manner shown in 
Fig. 46, using platinum electrodes and a sufficient number of 
dry cells or other source of current to give a suitable deflection 
of the galvanometer, we shall find that if we replace the solution 
by distilled water practically no current will be indicated by the 
galvanometer. We also find that if dry salt is substituted for 
the solution no current will pass. If now we pour distilled 
water on the salt while the latter is still in contact with the 
electrodes, a current begins to pass through the solution of salt 
which is quickly formed. 




&&- 



FlG. 46 

These results lead to the conclusion that neither pure water 
nor dry salt conducts the current appreciably compared with 
the solution formed from salt and water. All other soluble 
salts behave similarly. It is also easily discovered by experi- 
ment that dry (water-free) bases and acids are no better con- 
ductors than dry salts, although water solutions of acids and 
bases are good conductors. Water solutions of other substances 
than acids, bases, and salts do not conduct electricity any better 
than does pure water. 

The close connection which always exists between electrical 
conductivity and the ability of a mixture to undergo double 
decomposition is illustrated by the following experiment. Ferric 
sulfate and sodium, carbonate can be mixed dry without any 
apparent change; but let the mixture once be wet with water, 
immediately a violent evolution of gas follows and a red 
precipitate of ferric hydroxide appears. That the mixture of 
the dry substances is a non-conductor is shown by placing it 
in the dry beaker, Fig. 46. No current passes, but when water 
is added the substances dissolve, and the solution so formed 
conducts the current. At the same time the chemical reaction 



232 Introduction to General Chemistry 

begins vigorously. Since chemical reactivity and electrical con- 
ductivity seem therefore to go hand in hand, we shall next study 
the behavior of solutions of acids, bases, and salts when an electric 
current is passed through them. 

385. The Electrolysis of Solutions. — We have already learned 
that the electrolysis of concentrated hydrochloric acid sets free 
hydrogen and chlorine (43), and that the electrolysis of common 
salt (238) yields these same gases and in addition forms sodium 
hydroxide. In the case of hydrochloric acid, electrolysis simply 
causes the separation of the two constituents, 

HC1-»H+C1. 

On being set free the atoms of the two elements each form 
diatomic molecules, thus, 

2H^H 2 , and 2C1->C1 2 . 

These last reactions are doubtless secondary; and for the sake 
of brevity, in the examples of electrolysis that follow, reactions 
of this kind will be indicated by separate equations without 
further comment. 

In the case of the electrolysis of common salt it seems pos- 
sible, as already explained (238), that the first change is a 
separation into sodium and chlorine, thus: 

NaCl->Na+Cl, 
2C1->C1 2 . 

The sodium then reacts with the water present to form sodium 
hydroxide and hydrogen: 

Na+H 2 0->NaOH+H, 
2 H->H 2 . 

Whether this is the only possible explanation of the way the 
changes take place can best be discussed later; but it can be 
pointed out here that the foregoing equation would demand that 
the sodium hydroxide should be formed at the electrode at which 
the hydrogen is given off; and this is, in fact, the case. When 
hydrogen is set free in the electrolysis of any substance it always 



The Ionic Hypothesis 233 

appears at the negative electrode or cathode, while chlorine is 
liberated only at the positive electrode or anode. 

386. The Electrolysis of Copper Salts. — If a solution of 
cupric chloride, CuCl 2 , is electrolyzed between platinum poles 
or electrodes, copper is deposited on the negative pole and 
chlorine gas is set free at the positive pole. Here again, as in 
the case of hydrochloric acid, we have the simplest possible 
kind of a change, as represented by the following equation: 

CuCl 2 ->Cu+2Cl, 
2C1->C1 2 . 

If copper sulfate, CuS0 4 , is electrolyzed, a plating of metallic 
copper is again formed on the negative electrode, while from the 
positive electrode oxygen gas is given off. Examination of 
the products after electrolysis shows that sulfuric acid is con- 
tained in the solution surrounding the positive electrode. In 
fact, if the electrolysis is continued until all the copper in the 
original solution is deposited, all the sulfate radical of the 
original copper sulfate is changed into sulfuric acid, and this 
acid is contained in the part of the solution surrounding the 
positive electrode. The formation of sulfuric acid and oxygen 
may be explained by supposing the copper sulfate to be separated 
by the electric current into copper and sulfate radical, S0 4 , and 
that the latter reacts with water to form sulfuric acid and oxygen : 

S0 4 +H 2 0h>H 2 S0 4 +0, 

20->0 2 . 

387. The Electrolysis of Silver Nitrate. — If an electric cur- 
rent is passed through a solution of silver nitrate, AgN0 3 , silver is 
deposited on the negative electrode and oxygen and nitric acid 
appear at the positive electrode. Probably silver nitrate is 
first separated into silver and nitrate radical, N0 3 ; the latter 
then reacts with water to form nitric acid and oxygen: 

2N0 3 +H 2 0->2HN0 3 +0, 
20->0 2 . 

388. Summary. — In Table XIV we have summarized the 
results just discussed, leaving out of consideration the secondary 



234 



Introduction to General Chemistry 



changes that often take place between the substance set free and 
the water. We see that the parts into which a substance is separated 
by electrolysis are the same as those which change partners in double 
decomposition reactions. 

TABLE XIV 
Immediate Products of Electrolysis 



Liberated at 


Original 


Liberated at 


Negative Plate 


Substance 


Positive Plate 


H 


HC1 


CI 


Na 


NaCl 


CI 


Cu 


CuCl 2 


2C1 


Cu 


C11SO4 


so 4 


Ag 


AgN0 3 


N0 3 



389. The Terms Used in Electrolysis. — The phenomena of 
electrolysis were very carefully studied about 1834 by Faraday, 
who, as we shall see, discovered many important facts. It was 
Faraday also who invented the terms electrolysis, electrolyze, 
electrode, anode, and cathode. He called the solution the 
electrolyte, but at present we use this term to mean the dissolved 
substance. That part of the electrolyte which during electrolysis 
is deposited or set free at the anode or positive electrode he 
called the anion. The other part, which goes to the cathode, he 
called the cathion. Frequently he had occasion to speak of 
anions and cathions together, and then he referred to them as 
the ions of the electrolyte. For example, the ions of copper 
sulfate are said to be copper and sulfate radical, because in 
electrolysis copper passes to and is deposited on one electrode, 
while the sulfate radical goes to the other. Of course at some 
time or other the radicals or the partners of the original elec- 
trolyte must have broken apart, otherwise they could not have 
arrived at poles distant from each other. 

390. Hydrogen and Metals as Cathions. — We may next con- 
sider how the composition of the ions of a substance can be 
discovered. 

In the case of such a simple substance as HC1 it is obvious 
that the ions are hydrogen and chlorine, hydrogen being the 
cathion and chlorine the anion. Since all acids upon electrolysis 



The Ionic Hypothesis* 235 

give off hydrogen at the cathode, we may conclude that hydrogen 
is the cathion of all acids. 

When salts are electrolyzed the metal is either deposited in 
metallic form on the cathode, as in the case of copper and silver 
salts, or it collects in the solution about the cathode in the form 
of hydroxide, as when common salt is used. These facts lead 
to the conclusion that the basic or metallic elements of salts are 
cathions. 

391. Acid Radicals as Anions.— On the other hand, the acid 
elements or radicals of acids and salts accumulate at the anode 
and are either given off as free elements, as in the case of chlorine, 
bromine, and iodine, or they react with water to form acids and 
oxygen, as in the case of sulfate and nitrate radicals. 

392. Ions and Chemical Reactions. — It would thus appear, 
from what has just been stated, that the ions of an acid or salt 
are the same as the two parts of which its chemical reactions 
show it to be composed. It may be added that there is good 
reason for thinking that the same statement also applies to 
bases. The cathion of a base is usually a metal; the anion 
is the hydroxyl radical. 

393. Positive and Negative Ions. — The cathode is the electro- 
negative electrode; to it go the cathions. Since it is well known 
that a negatively charged body repels another negatively 
charged body and attracts one which is positively charged, it is 
not unreasonable to attribute the movement of ions to electrical 
attraction, and to conclude that cathions are electro positively 
charged. Since the anode is electropositive, the anions are 
thought to be charged electronegatively. It is customary to call 
cathions positive ions, and anions negative ions. 

394. The Cause of the Union of Ions. — Attention has been 
called to the fact that in reactions in solution basic or metallic 
radicals unite only with acidic or non-metallic radicals (383), 
and that unions of basic radicals with one another never occur; 
that is, double decompositions in water solutions never give 
compounds such as KNa and C1S0 4 . We are now in position to 
explain these important facts. We have learned that the radicals 
of acids, bases, and salts are identical with their ions, and that 



236 Introduction to General Chemistry 

the ions are probably electrically charged, the basic or metallic 
ones being positively, the acidic or non-metallic negatively, 
charged. We can therefore summarize by stating that in 
reactions only ions of unlike electric charges unite with one another. 
The reason for this is that ions with unlike electric charges 
attract each other, and that those with like charges repel each 
other. The chemical union of ions is an electrical phenomenon 
and is due to the attractive force of unlike electric charges carried 
by the ions. 

395. Colors of Ions. — The student has doubtless already 
observed that, although most salts and their solutions are color- 
less, a considerable number are colored. A little investigation 
will show that very dilute solutions of equal concentration of 
all cupric salts of colorless acids are of the same shade and 
intensity of blue color. This fact leads us to believe that the 
blue color is due to copper ions, which is the only substance which 
all the solutions have in common. Moreover, we find that the 
colors of all dilute solutions of colored acids, bases, and salts 
can be ascribed to the colors of their ions. 

If the dilute solution of any acid, base, or salt is colorless, 
like pure water, we may conclude that its positive and negative 
ions are both colorless. If a dilute colored solution of an elec- 
trolyte has one colorless ion we conclude that the observed 
color is that of the other ion. Thus we find that all dilute 
ferrous solutions (173, 331) are pale green and conclude that 
ferrous ion is pale green. Manganous salts (342) (for example, 
MnCl 2 and MnS0 4 ) give pale pink solutions, therefore positive 
Mn ion is pale pink. On the other hand, dilute solutions of all 
permanganates (343) are, like KMn0 4 , deep purple, and we 
conclude that negative Mn0 4 ion is purple. Similar reasoning 
leads us to conclude that negative Cr0 4 ion is yellow (345) and 
negative Cr 2 7 ion is orange (345), while positive Cr ion is violet 
(344). The color of a dilute solution is usually an indication of 
the nature of one of its ions. 

396. Colors of Molecules. Dissociation of Molecules into 
Free Ions. — Although dilute solutions of all cupric salts are 
blue the solid salts and also their concentrated solutions are in 



The Ionic Hypothesis 237 

several cases of a different color. Thus cupric chloride, CuCl 2 , 
in solid form and in concentrated solution is green, and cupric 
bromide, CuBr 2 , which forms almost black crystals, gives a 
concentrated solution which is dark brown; but if this brown 
solution is sufficiently diluted the color gradually changes to 
blue, finally reaching the same shade of color as that of any 
other equally dilute cupric solution. A simple explanation 
of these color changes is found in the assumption that the dark 
brown color is that of the molecules, CuBr 2 , while the blue color 
is due to Cu ions. From the fact that many dilute solutions 
of bromides are colorless we conclude that Br ions are colorless. 
By following up this idea we are led to a very remarkable con- 
clusion, namely, that molecules of CuBr 2 exist only in the solid 
state and in concentrated solutions but not to an appreciable 
extent in very dilute solutions. This is accounted for if we 
assume as the concentrated solution is diluted molecules gradually 
split up or dissociate into ions, thus : 

CuBr 2 ^Cu-f-2Br, 

so that in a dilute solution the substance exists largely as free Cu 
and Br ions. If we evaporate the dilute blue solution we again 
obtain a brown concentrated solution and finally brown crystals. 
We must therefore assume that the change is a reversible one, the 
ions reuniting to form molecules as the solution is evaporated. 
Further evidence of the existence of free ions is afforded by the 
experiments next to be considered. 



Ss_ 



Fig. 47 

397. The Migration of Ions. — Let us take advantage of the 
color of ions to discover their behavior during the process of 
electrolysis. In the U-tube, Fig. 47, we may put a solution of 
a colored electrolyte in the lower layer and colorless electrolytes 



238 Introduction to General Chemistry 

in the layers next to the electrodes. As colored electrolytes we 
may use copper nitrate or potassium permanganate. When the 
current is turned on, the boundary of each, colored electrolyte 
slowly moves up into one of the colorless layers above it, just 
as we would expect if the colored materal is the free ion which 
carries a charge of electricity. Thus positive copper ion migrates 
toward the negative electrode, and negative permanganate ion 
migrates toward the positive electrode. We can carry out an 
experiment with a mixture of these two colored salts in the lower 
layer. The purple layer now shows on the side of the positive 
electrode, and the blue layer shows on the side of the negative 
electrode just as before. Thus we find that each ion migrates 
just as though the other were not there; and this, in fact, is 
just what we should expect if a dilute solution contains free 
ions formed by dissociation of its molecules. 

398. The Mechanism of Electrolysis. — We can now make a 
fairly complete picture of the mechanism of the conduction of 
the current through a solution and of the accompanying elec- 
trolysis. We shall assume that in the dilute solution the dis- 
solved substance is partially dissociated into positive and 

negative ions. Fig. 48 
represents diagrammati- 
cally such a solution 
in which the two elec- 
trodes are dipped, con- 
nected with a pair of 
dry cells. The cells 
charge the electrodes, 
one positively, the other negatively. The influence of these 
charges is felt by the ions in the solution. The positive ions 
are attracted by the negative electrode and repelled by the 
positive electrode and in consequence migrate toward the 
former. The negative ions move in the opposite direction for 
similar reasons. 

When ions come into contact with the electrodes of unlike 
sign they give up their charges to the electrodes. This tends 
to discharge the latter, but the battery keeps them charged by 



=> 6 e s e 

©«© e 

sF3 ej- e 

i%fe 9 e 

e©©e e 



L_g_. 



f^ 



Fig. 48 



The Ionic Hypothesis 



2 39 



continuously sending a current of electricity through the wires. 
A more detailed description of the mechanism of electrolysis, 
in terms of the newer views of the nature of electricity, will be 
given in chapter xx. 

399. Faraday's Laws of Electrolysis. — As the result of care- 
ful experimental investigation of the quantities of substances 
liberated during electrolysis, Faraday arrived at the following 
conclusions : 




«-^£ 



r 



Fig. 49 

1. The amount of a given substance, say hydrogen, set free by 
electrolysis is directly proportional to the quantity of electricity 
which is passed through the solution. 

2. The amount of a substance, hydrogen for example, which 
is liberated by a fixed quantity of electricity is the same, whatever 
the nature of the solution electrolyzed, provided that this substance 
and nothing else is liberated at the given electrode. These two 
statements are known as Faraday's Laws of Electrolysis. 

400. Two Electrical Units. — To understand these laws fully 
we must review briefly some fundamental facts about the elec- 
trical current so that we can appreciate what is meant by 
quantity of electricity. In the first place we know that if a 
current passes through a wire there is produced around the wire 
a magnetic field. If we attach a thread to the middle of a 
magnetized steel needle and suspend the latter above and 
parallel to a wire, then as soon as we cause a current of elec- 
tricity to pass through the wire the needle sets itself at an angle 
to the former, Fig. 49. The greater the angle between the 



240 



Introduction to General Chemistry 




Fig. 50 



needle and the wire the stronger is the magnetic field, and the 
stronger the current is said to be. It is on this principle that 
instruments are built to measure current strength. Of course, 
to measure anything we must first adopt some fundamental 
unit by comparison with which the measurement can be made. 
This was done in the case of the electric current on the basis of 
the strength of the magnetic field about a conductor, and this 
unit was called the ampere. The ammeter 
(Fig. 50) allows us to read, from the posi- 
tion occupied by the needle on its scale, just 
how many amperes of current are passing. 
The amperage tells us the strength of 
the current, but we must also know the 
time during which the current is allowed to 
pass if we are to know the amount of elec- 
tricity delivered at the terminals of the con- 
ductor, say at two electrodes. If a current of one ampere is 
allowed to flow one second it is said to deliver a unit quantity of 
electricity, and this unit is called the coulomb. 

401. Illustration of Faraday's Laws. — The following facts 
will serve to illustrate the meaning of Faraday's laws. By the 
electrolysis of dilute acids hydrogen is set free at the negative 
electrode. In all cases the passage of 96,500 coulombs of 
electricity is required for the liberation of one gram of hydrogen. 
Since a current of one ampere delivers one coulomb per second, 
96,500 coulombs will be given by a current of one ampere in 
96,500 seconds, or 26.8 hours. A current of two amperes for 
the same length of time will liberate 2 g. of hydrogen, or one 
gram molecular weight of hydrogen (H 2 ) , which as we know has 
a volume of 22.4 liters at o° and 76 cm. 

402. Discussion. — It is not surprising that if a one-ampere 
current will liberate 1 g. of hydrogen in 26. 8 hours, a two-ampere 
current will liberate 2 g. of hydrogen in the same time, for this 
is the type of regularity which we have become accustomed to 
expect in nature. It is surprising, however, that the same 
amount of hydrogen is liberated by the same amount of elec- 
tricity from a solution of any dilute acid, and the fact that this 



The Ionic Hypothesis 



241 



is so must reflect some regularity in the phenomena of elec- 
trolysis, the cause of which we have still to discover. 

403. Faraday's Laws of Electrochemical Equivalents. — Let 
us now turn to cases of the liberation by electrolysis of elements 
other than hydrogen. Very careful experimentation has shown 
that by the passage of 96,500 coulombs of electricity through 
various solutions certain weights of elements are set free. These 
are given in Table XV. This table shows a most striking regu- 

TABLE XV 

Electochemical Equivalents 



Element 


Gram Atomic 
Weight 


Valence *<£■£>£** 


Gram Atomic 
Weight-e- Valence 


Hydrogen 

Silver 


1 
108 

64 

66 
208 

56 

27 

35-5 

80 

16 


I 
I 

2 
2 
2 
2 

3 

I 
I 
2 


I 

I08 

32 

33 
104 

28 

9 
35-5 
80 

8 


I 
108 


Copper 


32 

33 

104 


Zinc 


Lead 


Iron 


28 


Aluminum 

Chlorine 

Bromine 

Oxvgen 


9 
35-5 
80 

8 









larity: The weight of an element liberated in electrolysis by the 
passage of 96,500 coulombs of electricity is equal to the gram atomic 
weight of that element divided by its valence (col. 5). This weight 
is called the electrochemical equivalent of a given element or, 
more briefly, its equivalent weight. The electrochemical equiva- 
lents of the various elements are seen to be proportional to the 
weights of these elements which unite chemically with one 
another when union is possible; for example, 1 g. of hydrogen, 
104 g. of lead, or 9 g. of aluminum unite with 35 . 5 g. of chlorine, 
or 8 g. of oxygen. The discovery of facts such as those given 
in the table was made by Faraday, who stated his conclusion as 
the Law of Electrochemical Equivalents : The amounts of different 
substances liberated by the same quantity of electricity are propor- 
tional to their equivalent weights. 

404. The Electric Charges of Ions. — The facts covered by 
Faraday's laws allow us to draw some interesting and significant 



242 Introduction to General Chemistry 

conclusions regarding the quantities of electricity composing 
the charges on single ions. If the 96,500 coulombs of electricity 
supplied at the negative electrode to release one gram of hydro- 
gen ion are used simply to neutralize the charge on one gram of 
that ion, we may conclude at once that the charge carried by 
the one gram of hydrogen ion is not only opposite in sign but 
equal in amount to the electricity required. In general then 
one formula weight of a univalent ion must carry a total charge 
equal to 96,500 coulombs. One formula weight of an ion of any 
valence will carry one, two, three, four, etc., times this charge, 
according to whether its valence is one, two, three, or four, etc. 
If we assume that one formula' weight of any one ion represents 
the same number of free ions as a formula weight of any other 
ion (and this is in strict accord with our accepted definition of 
the term formula weight) , we come at once to the conclusion that 
all univalent ions carry equal charges. We call this a unit charge ; 
each bivalent ion carries two unit charges, each trivalent ion carries 
three unit charges, etc. In writing the symbols or formulae of 
free ions it is customary to add one or more -f- or — signs to 
indicate the number of positive or negative unit electric charges 
carried by the ion, for example, H + , Cu ++ , Al +++ , Cl~, 

so 4 ~, po 4 — . 

405. Equilibrium between Molecules and Ions. — The facts 
already studied (396), together with a great volume of other 
evidence which we shall take up in turn, led the Swedish chemist 
Svante Arrhenius to the conclusion that in concentrated solu- 
tions of acids, bases, and salts a considerable part of the dissolved 
substance is present as molecules; but that as the solution is 
diluted, more and more of the molecules dissociate into free ions, 
until in very dilute solutions (at least in many cases) the disso- 
ciation is nearly complete. On the other hand, when a dilute 
solution is evaporated the ions undoubtedly gradually unite 
to form molecules, until, when complete dryness is reached, only 
molecules are present. In any given solution a state of equi- 
librium exists between molecules and ions, as represented in the 
case of common salt by the equation 

NaCl^Na++Cl-. 



The Ionic Hypothesis 



243 



At a definite concentration a definite proportion of the salt will 
be present as ions; this proportion we call the fraction ionized 
or the degree of ionization. We shall next take up the important 
problem of determining the fraction ionized for any solution of 
an electrolyte. Since we believe that the current in a solution 
is carried by the ions present, the ability of a solution to conduct 
a current, or, briefly, its conductivity, must be an indication of 
the extent to which its molecules are dissociated into ions. 

406. Effect of Dilution on Conductivity. — We have already 
learned from the color changes produced by diluting solutions 
that ionization is promoted by dilution. Let us now consider 
the question, What influence, if any, does the volume of water 
in which a given quantity of an acid, base, or salt is dissolved 




Se-rSf-^ 



Fig. 51 

have on its electrical conductivity? We may study this ques- 
tion experimentally by means of the apparatus shown in Fig. 5 1 . 
The rectangular glass vessel of about 1 liter capacity is provided 
with two large copper electrodes, as shown in the figure. The 
vessel is first filled about three-fourths full of distilled water, 
and the electrical connections are made. No appreciable current 
passes. Next about 200 c.c. of concentrated hydrochloric acid 
are introduced below the water, without mixing, in such a way 
as to form a separate layer. This may be done by the use of a 
dropping funnel, the stem of which reaches the bottom of the 
vessel. 

The vessel now contains two distinct layers — a lower layer 
of concentrated hydrochloric acid and an upper layer of water. 
The galvanometer indicates that a considerable current is 
passing, and we conclude that this is all passing through the 
acid in the lower layer and not through the upper water layer. 



244 Introduction to General Chemistry 

If next we mix the acid and water thoroughly and so dissolve 
the acid in a much larger volume of water, we note that a large 
increase in the current takes place. This leads us to conclude 
that the conductivity of the hydrochloric acid present is greater 
when it is dissolved in the larger volume of water. We may now 
ask, however, whether there is a limit to the increase in conduc- 
tivity when a given amount of acid is dissolved in larger and 
larger volumes of water, the conductivity being measured under 
such conditions that the solution is all contained between parallel 
plates at a fixed distance apart. 

If, in the experiment described, the vessel were much deeper, 
but otherwise the same, and the electrodes extended all the way 
up the sides as before, it would be found that a given amount 
of hydrochloric acid diluted with double the amount of water 
used in the first experiment would show appreciably greater 
conductivity than in the first case. Or if the acid were diluted 
with three, or four, or still more times as much water, greater 
and greater conductivity would have been observed; but with 
increasing dilution the increase in conductivity would become 
smaller each time more water was added, so that finally a maxi- 
mum conductivity would be reached. Beyond this limit further 
dilution would cause no increase in conductivity. 

These same experiments could be repeated with many other 
electrolytes with similar results. 

407. Definition of Molecular Conductivity.— If one formula 
weight (called also one gram molecular weight) of an acid, base, 
or salt is contained in a solution which is wholly included between 
two parallel electrodes 1 cm. apart, we call the electrical con- 
ductivity of this solution its molecular conductivity. To find 
the molecular conductivity experimentally we measure its 
electrical resistance in ohms. The reciprocal of the resistance 
so found is by definition the molecular conductivity. The con- 
clusions of the paragraph on the effect of dilution on conductivity 
may now be summarized as follows: The molecular conductivity 
of every electrolyte increases as its solution is diluted and finally 
attains a maximum which has a definite numerical value for each 
substance (the temperature being fixed). Table XVI shows the 



The Ionic Hypothesis 245 

change of molecular conductivity of hydrochloric acid as the 

volume in which one formula weight of acid is contained is 

increased. 

TABLE XVI 

The Molecular Conductivity of Hydrochloric 
Acid at i8 c (No yes and Cooper) 

Volume of Solution Molecular 

in Liters Conductivity 

IO 351 

12.5 353 

100 368 

500 373 

2,000 375 

Maximum 379 

408. Determination of the Degree of Ionization — When a 
solution of a substance is so dilute that it has its maximum molec- 
ular conductivity, it is assumed that all of its molecules have 
dissociated into ions. In a more concentrated solution, for which 
the molecular conductivity is less than the maximum, the frac- 
tion which its observed molecular conductivity forms of its 
maximum molecular conductivity is consequently equal to the 
fraction which the number of ions present in that particular 
solution form of the total number of ions in the completely 
dissociated (completely ionized) solution of the same quantity 
of that substance. This fraction is therefore the fraction ionized 
or the degree of ionization. Thus for decinormal hydrochloric 
acid the degree of ionization is 351-^-379 = 93 per cent. 

This method of determining the degree of ionization was 
proposed by Arrhenius in the year 1887. His reasoning ran 
thus: The passage of a current through a solution is accom- 
plished by the migration of positive ions in one direction and 
negative ions in the other. These transport electricity through 
the solution between the electrodes. Since the molecular con- 
ductivity of a substance is the measure of its rate of transporting 
electricity, it is plain that the molecular conductivity will depend 
on the number of ions present, the charge on each, and the velocity 
of migration. Now under the conditions used in measuring 
resistance, and therefore also of measuring molecular con- 
ductivity, the velocity of migration of its ions will be the same 



246 Introduction to General Chemistry 

for all concentrations of solutions of a given substance (except 
for very concentrated solutions). The charges of the individual 
ions of a given substance are also the same, whether the solution 
is dilute or concentrated. Therefore the molecular conductivities 
of solutions of a given substance are directly proportional to the 
numbers of ions present. Consequently the ratio of the molecu- 
lar conductivity for a given concentration to the maximum mo- 
lecular conductivity for this substance is the fraction ionized, since 
it is assumed that a very dilute solution having maximum molec- 
ular conductivity is completely ionized. 

409. Results of Determination. — The degree of ionization of 
some common electrolytes is shown in Table XVII. 

410. Discussion of Table XVII. — A study of the table leads 
to the very important generalization that in solutions of most 
salts a large percentage of the substance is in the form of ions; in 
consequence, we say that such solutions are highly ionized. It 
also appears that dilute solutions of hydrochloric and nitric acid 
are even more highly ionized than salt solutions of like concentra- 
tion. On the other hand, decinormal acetic acid is only 1 . 3 per 
cent ionized, while the degree of ionization of decinormal carbonic 
acid is very much less, namely 0.17 per cent. In general, the 
extent to which acids are ionized, in solutions of equal concen- 
tration, varies enormously. Bases also differ greatly in their de- 
grees of ionization. For example, decinormal sodium hydroxide 
is ionized 90 per cent, while the same concentration of ammonium 
hydroxide is only 1 . 3 per cent ionized. We have already learned 
that every substance is more highly ionized in dilute tfyan in 
more concentrated solutions. The percentage of ionization of 
a substance as shown in the table applies only to the indicated 
concentration and temperature. 

411. Resume of the Ionic Hypothesis. — We have already 
developed enough of the ionic hypothesis to go far into the under- 
standing of double decomposition reactions. Let us therefore 
review in brief the ideas already brought out, and then, after 
a short critical survey of the fundamental assumptions, proceed 
in chapters xviii and xix to the application of the hypothesis to 
practical examples. 



The Ionic Hypothesis 



247 



According to the ionic hypothesis, as soon as an acid base or 
salt is dissolved in water it is immediately dissociated to some 
extent into ions which prove to be the parts of those substances 
which we have found active in double decomposition. The basic 

TABLE XVII 

Values of the Degree of Ionization of Some Common Electrolytes in 
Water Solution at i8° 

(Degree of ionization at the normality indicated at. the head of the column) 



Salts : 

Nad 

KC1 

KBr 

KI 

NaN0 3 

KNO3 

AgN0 3 

KCIO3 

BaCl 2 

CaCl 2 

MgCl 2 

PbCl 2 

Sr(N0 3 ) 2 

Ba(N0 3 ) 2 

K 2 S0 4 

Ag 2 S0 4 

MgS0 4 

ZnS0 4 

CuS0 4 

Bases: 

NaOH 

Ba(OH) 2 

NH 4 OH 

Acids: 

hq...... 

HN0 3 

HC 2 H 3 2 

H 3 P0 4 = H++H 2 P0 4 - 
H 2 S0 4 = 2 H++S0 4 — 
H 2 C0 3 = H++HC0 3 - 



0.94 
94 

94 
94 
93 
94 
93 
93 
88 



93 

04 

97 

97 

042 

60 

64 

005 



46 

46 

93 
86 
017 



94 
94 
020 
36 

38 
002 



45 
4i 
40 

90 
81 
013 

93 
93 
013 

29 
3i 
0017 



0.74 
74 



5i 



59 



part carries a positive charge and the acidic part a negative 
charge. The charge carried by any univalent ion is called a 
unit charge; ions having greater valence carry as many unit 
charges as they have valence. Since the solution of any elec- 
trolyte is always electrically neutral, the total quantity of posi- 
tive electricity carried by the positive ions equals the total 



248 Introduction to General Chemistry 

quantity of negative electricity carried by the negative ions. 
The more dilute the solution the greater is the proportion of the 
electrolyte transformed into ions, or, in other words, the greater 
is its degree of ionization. 

When we put two electrodes, connected with dry cells or 
other source of current, into a solution of an electrolyte, the 
current is found to flow in the outside circuit, because of the 
discharging of the charge on the electrodes, due to the arrival 
at their surface of oppositely charged ions from the solution. 
The ions in the solution move up to the plates because each ion 
is attracted toward one plate and repelled from the other, owing 
to the fact that it also carries a charge of electricity. The sign 
of the charge on the ion determines the direction of the latter's 
movement. On coming in contact with the electrodes the ions 
become discharged by having their charges electrically neutral- 
ized by equal amounts of the opposite kind of electricity furnished 
by the electrode. Metallic ions after discharge are either de- 
posited as metallic platings on the cathode or react with water 
to form hydroxides and hydrogen. Non-metallic ions such as 
hydrogen, oxygen, and chlorine are released as single atoms which 
then unite to form diatomic molecules of the gases H 2 , 2 , Cl 2 , 
etc. Nitrate and sulfate ions never remain discharged at the 
electrodes, but instead we find there the products of their 
reactions with water — nitric acid and oxygen, and sulfuric acid 
and oxygen, respectively. 

If a given quantity of electrolyte is kept between plates 
which are parallel to each other and carry a constant charge per 
unit area, at different dilutions the conductivity of this elec- 
trolyte will vary in proportion to the number of free ions present. 
As a consequence the proportion of an electrolyte which has 
been transformed into ions at any dilution can be determined 
by dividing the conductivity at the dilution in question by the 
maximum conductivity found after continuous dilution of this 
same quantity of electrolyte, provided the two measurements 
are made in the manner described (408). Values so determined 
show us that as a rule all salts are highly ionized substances, 
but acids and bases have very different degrees of ionization, 



The Ionic Hypothesis 249 

some being even more highly ionized than salts, but others 
being very little ionized indeed. 

412. Criticisms of the Ionic Hypothesis. — The idea that ions 
exist in solution as independent chemical substances has come 
to be known as the Ionic Hypothesis. It will be surprising if 
the student who learns of this hypothesis for the first time and 
thinks critically about the matter is not ready to offer at once 
several good reasons for doubting the truth of the conclusions. 
In the first place, the hypothesis seems to assume that in a solu- 
tion of common salt, for example, a large part of its elements 
are present in a free state. Now the student who knows any- 
thing of the properties of metallic sodium and of chlorine gas 
will find it hard to believe that either of these elements can be 
present in a salt solution; because sodium reacts violently with 
water, forming sodium hydroxide and hydrogen, and chlorine 
has a horrible smell and a yellow color. Plainly there is some- 
thing incompatible with the obvious facts in the statement that a 
solution of salt contains free sodium and free chlorine. 

A closer study of the hypothesis shows, however, that it is 
not assumed that the elements sodium and chlorine are present 
as ordinary atoms, for each atom is said to be electrically charged. 
Those who uphold the hypothesis will point out that a charged 
brass ball has very different properties from the same ball if 
uncharged. True, say the critics, but even a charged brass ball 
is still a brass ball; to which the opponents reply that the 
quantity of electricity on the ball is a matter of enormous 
importance. 

If then the ions are so highly charged, why do the positive 
ions not unite with and so electrically discharge the negative 
ions, since the solution is a conductor? It may be said in reply 
that it is assumed that ions of unlike sign are constantly uniting, 
at a rate just equal to the rate of dissociation, with the result 
that a state of equilibrium is produced. 

In spite of the foregoing criticisms and many others the ionic 
hypothesis with all its apparent inconsistencies has proved itself 
highly useful in explaining and correlating many facts and 
phenomena. 



250 Introduction to General Chemistry 

Before passing final judgment on this remarkable hypothesis 
it will be better to consider its further applications and then, 
in chapter xx, to take up the matter again in the light of newer 
discoveries, which have led to essential modifications of the 
views as originally proposed by Arrhenius. 

Finally it may be urged that Arrhenius himself was not 
certain of the truth of his theory until he became acquainted 
with the wonderful work of Van't Hoff on the so-called osmotic 
pressures of dissolved substance (chap, xxvii). This work 
will be discussed as soon as we have progressed far enough to 
understand and interpret the experiments which we must then 
study. 



CHAPTER XVIII 
APPLICATIONS OF THE IONIC HYPOTHESIS 

413. Double Decomposition. — In the foregoing chapter it 
was pointed out that the probable cause of the union of two 
unlike ions is the attraction of their unlike electric charges. In 
general, every kind of positive ion can unite with any kind of 
negative ion. Therefore, if any two electrolytes (provided they 
have no ion in common) are mixed in solution, at least some 
double decomposition must take place, simply because new 
combinations of positive and negative ions are made possible. 
Let us first consider the important case in which the two starting 
materials, as well as the two products of the reaction, are easily 
soluble and highly ionized. 

414. Class I. Equilibrium between Four Easily Soluble and 
Highly Ionized Electrolytes.— If dilute solutions of two imaginary 
electrolytes A B and CD, which ionize thus 

AB±*A++B-, I 

CD^C++D~, II 

are mixed, we may predict, without knowing anything more 
about these substances, that the following reactions are possible, 

A++D-±+AD t III 

C++B~^CB, IV 

and that a double decomposition reaction, 
- AB+CD±^AD+CB, 

will take place to a greater or less extent. Since all four of the 
substances AB, CD, AD, and CB are assumed to be highly 
ionized, it is plain that the mixed solution will contain chiefly the 
four kinds of ions, A + , B~ , C + , and D~ , and relatively few 
molecules. Since each of the four kinds of molecules present 
must be in equilibrium with its own two kinds of ions, the four 

251 



252 Introduction to General Chemistry 

equilibrium reactions (marked I, II, III, and IV) must be 
interrelated in the manner shown by the following arrangement 
of equations I, II, III, and IV: 

III IV 

AB^A++B~ I 

CD*=>D-+C+ II 

it It 
AD CB 

Equations I and II read horizontally, while III and IV read 

vertically. We may call this the compound equation of the 

reaction 

AB+CD±^AD+CB. 

The compound equation shows that the four molecular sub- 
stances are in equilibrium with each other because each molec- 
ular substance is in direct equilibrium with its own pair of ions- 
Now if all of the four molecular substances are assumed to have 
exactly equal tendencies to ionize, then we must conclude that 
for the condition of equilibrium equal numbers of the four kinds 
of molecules will be present, if we have taken equivalent amounts 
of substances. We may summarize Class I as follows: If both 
starting substances and both products of a double decomposition 
reaction" AB+CD^ AD+CB are easily soluble and highly 
and equally ionizable, an equilibrium mixture will result in 
which (1) most of the dissolved material is present as free ions, 
(2) little of the material is present as molecules, and (3) if 
equivalent amounts are taken the four kinds of molecules are 
present in equal numbers. 

415. An Example of Class I. The Reaction between Ferric 
Chloride and Ammonium Sulfocyanate. — The reaction . 
FeCl 3 +3NH 4 NCS^Fe(NCS) 3 +3NH 4 Cl, 

studied earlier (280), is a good illustration of Class I, since all 
four salts are easily soluble and highly ionized. It was shown 
by experiment that this reaction does not take place completely, 
but that it reaches equilibrium while there is still much of the 
material not converted into ferric sulfocyanate and ammonium 
chloride. 




Applications of the Ionic Hypothesis 253 

416. Other ■ Examples of Class I.— Numerous additional 
examples of Class I might be given. The following will serve 
as illustrations: 

KCl+NaN0 3 ^KN0 3 +NaCl, 

Na 2 S0 4 + 2 KN0 3 ^K 2 S0 4 + 2 NaN0 3 , 

K 2 C0 3 +Na 2 S0 4 ±5K 2 S0 4 +Na 2 C0 3j 

NaN0 3 +HCl±5NaCl+HN0 3 . 

In each case the mixed solution contains largely the four kinds 
of ions, together with small proportions of the four kinds of 
molecules in approximately equivalent amounts. In Class I 
the two substances taken react only partially, and there- 
fore the reaction is always incomplete. 

417. A Graphic Method of Representing Degrees of 
Ionization. — An acid, base, or salt, not in solution, exists 
wholly in the form of molecules (no ions are present). 
We may represent such an un-ionized substance by a cross- 
hatched circle, Fig. 52. When this substance, whose formula we 
may call AB, is dissolved in water it will partially ionize, thus: 

AB^A++B~. 

This condition is represented by Fig. 53. Let us suppose that 
the solution is 80 per cent ionized ; then 20 per cent is present as 
un-ionized molecules. In Fig. 53 the left-hand circle has a 

cross-hatched sector which is 

''~^\AB 4iiiiM + ^^ r j ust 2o p er cent °^ ^ e area 

^F ' ^ \ >g 1 y^^N of the whole circle. This will 

^ _ ' ^ujiiiip^ represent the fact that 20 

F IG . 53 per cent of the substance is 

present as un-ionized mole- 
cules. The middle circle, of which 80 per cent is shaded with 
vertical lines, will represent the fact that 80 per cent of the 
total A radical is in the form of free positive ions. In similar 
fashion the left-hand circle shows that 80 per cent of the 
total B radical is in the form of free negative ions. Further- 
more, if we take the area of the circle, Fig. 52, as propor- 
tional to the whole number of molecules in one formula weight 
of the substance before it is dissolved, then the area of the 



254 



Introduction to General Chemistry 




CuSO* 



CuSO. 




so.-- 



cross-hatched sector of the left-hand circle of Fig. 53 will be pro- 
portional to the number of un-ionized molecules in one formula 
weight of the dissolved substance. Since each AB molecule, 

when it ionizes, gives one A + 
ion and one B~ ion, the areas 
of the shaded portions of the 
middle and right-hand circles 
will be directly proportional to 
the numbers of A + and B~ 
ions respectively. By means 
of a figure like Fig. 53 the 
relative concentrations of ions 
and molecules of a dissolved 
electrolyte can be seen at a 
glance. By way of further 
illustration the condition of 



- N CuSO. 




Fig. 54 



normal, one-tenth-normal, and one one-hundredth-normal copper 
sulfate solution is shown in Fig. 54. 

418. Graphic Representation of Class I. — Let us now turn 
to the graphic representation of a double decomposition reaction 
of the type just studied under Class I, where all four substances 
concerned are easily soluble and highly ionized. We again repre- 
sent the reaction by 

AB+CD^AD+CB. 

Figure 55 shows the condition of solutions of AB and CD 
before they are mixed, on the supposition that each is 85 per cent 
ionized in N/10 solution. When 
equal volumes of the two N/10 
solutions are mixed, the reaction 
represented by the compound 
equation 

AB^A++B~ 
CD^D-+C+ 
.It It 
AD CB 




Fig. 55 



takes place and very quickly reaches the condition of equilib- 
rium shown graphically in Fig. 56, in which the proportions of 




Applications of the Ionic Hypothesis 255 

molecules and ions have been calculated on the additional 
assumption that AD and CB both have the same tendency to 
ionize as have AB and CD (when each is separately dissolved 
in water). Comparison of Figs. 55 and 56 shows us that the 
areas representing the numbers 
of molecules and ions of the 
materials taken are not greatly 
changed as the result of the 
mixing. Consequently we say 
that the reaction is incomplete. 
All examples of Class I would be 
represented by similar graphs. 

419. A Second Type of 
Double Decomposition: Class 
II. — Class II will comprise Fig. 56 
double decomposition reactions 

in which two easily soluble and highly ionized substances give 
two easily soluble products, one of which is highly ionized, the 
other little ionized. The simplest example of Class II is found 
in a neutralization reaction such as 

HC1+ NaOH ±? NaCl+ H 2 0, 

since all the substances except the water are highly ionized. 

420. The Ionization of Water. — The ionization of water may 
be determined from conductivity measurements, for though it is 
a very much poorer conductor of the current than is a salt solu- 
tion, still, as we have already said, it conducts much better 
than glass or hard rubber. According to the ionic theory it is 

ionized thus: 

H 2 O^H++OH". 

In one liter of pure water there is present about one ten-millionth 
of a gram of ionic hydrogen and the equivalent amount of 
hydroxyl. 

If then we attempt to represent the proportion of ions in pure 
water by a graphic scheme, a single dot in the center of an other- 
wise empty circle would have too large an area to represent 
correctly the proportion of ions present if the rest of the circle 




256 Introduction to General Chemistry 

represented the molecules of water. In cases of this kind we 
shall use a single dotted radius to indicate that the number of 
ions is too small to be accurately represented. The graph of 

water will then be that shown in 

; { —\ { ---) Fig. 57. 

V_/ v . _,.' That there are so few ions 

FlG present in a liter of pure water 

means that the ionic equilibrium 
is established only when all but a minute fraction of the total 
material is in the form of water molecules. Accordingly, when 
hydroxyl and hydrogen ions are brought together in solution 
we must expect them to combine almost completely to form 
molecules. 

421. Neutralization. — If we mix equivalent amounts of 
solutions of HC1 and NaOH the resulting reaction may be 
represented as follows: 

HC1^H++C1- 
' NaOH^OH"+Na+ 

It It 

H 2 NaCl. 

The H + and OH~ ions present unite almost completely to form 

molecules of H 2 0. The removal of H + ions causes a shift of the 

reaction 

HC1^H++C1" 

to the right, and as the H + ions produced in this way are almost 
immediately taken up by new OH~ ions formed by a shift to the 
right of the reaction 

NaOH^OH-+Na+, 

the final result is the practically complete dissociation of both 
HC1 and NaOH molecules and therefore the disappearance of 
these substances. Molecules of H 2 0, once formed, dissociate 
very little into H + and OH~ ions, and so the final equilibrium 
solution will contain no more free H + and OH~ ions than an equal 
volume of pure water. The Na + and CI - ions unite partially to 

form molecules 

Na++Cl-^NaCl, 



Applications of the Ionic Hypothesis 



257 



but this reaction does not proceed far in dilute solution, as com- 
mon salt is a highly ionized substance. In fact, the solution 
resulting from the neutralization of HC1 by NaOH is exactly 
the same as, and differs in no way from, that made by dissolving 
common salt in water to produce a solution of equal concentra- 
tion. All the facts just stated are shown by a comparison of the 
two graphs, Figs. 58 and 59. Thus it can be seen (Fig. 59) that 
the areas representing the numbers of molecules of HC1 and 
NaOH respectively have been reduced to negligible dimensions; 
the same is also true of the areas for H + and OH~ ions. But 



,HCI 

\ 

1 ; 

/ 

NaOH 



■s. ' 





Fig. 58 



Fig. 59 



the Cl~ and Na + ion areas are not greatly changed in the second 
graph. Compared with these areas, that representing NaCl mole- 
cules is small. The circle representing the number of molecules 
of H 2 is completely shaded, thus showing that the yield of 
molecular H 2 is practically 100 per cent. 

422. The Simplified Equation of Neutralization. — To sum 
up the matter, it may be said that acids and bases neutralize 
one another because of the tendency of H + and OH~ ions to unite 
almost completely to form water. This almost complete union 
of H + and OH~ ions takes place because H 2 is but very slightly 
ionized. In a very dilute solution, where the acid and base 
taken are almost completely ionized at the instant of mixing, 
the principal change that takes place is the union of H + and 
OH~ ions to form H 2 molecules, since in the very dilute solu- 
tion the Na + and Cl~ ions remain largely uncombined. We may 



258 Introduction to General Chemistry 

therefore write as the simplified equation of neutralization in 
dilute solution 

H++0H-±5H 2 0. 

423. Experimental Confirmation of the Theory of Neutraliza- 
tion. — The process of neutralization can be followed experi- 
mentally with the help of an apparatus somewhat like that shown 
in Fig. 51 (406); but having a small electric lamp in the place 
of the galvanometer. The solution layers in the cell shown 
in the figure are made by first putting into the cell a layer of 
one-tenth-normal hydrochloric acid, and then allowing an equal 
layer of sodium hydroxide to run under this first layer by intro- 
ducing it at the bottom of the cell through a dropping funnel. 
As represented in the figure the two parallel electrodes are in 
contact with the two layers, which can be seen very nicely if a 
little litmus is put into the acid and base respectively before the 
layers are made. If now the key is closed the current flows 
through both layers, and the lamp glows. Hydrogen and sodium 
ions are arriving at one electrode, and chlorine and hydroxyl 
ions are arriving at the other. Of these ions the hydrogen and 
hydroxyl travel much more rapidly under the attraction from 
a given charge per unit area of the electrode, and so they are 
neutralizing their charges on the plates more quickly than are 
the other ions. As a result most of the current passing in the 
outside circuit is due to their discharge. If the two layers of 
acid and base are next mixed, the lamp no longer glows. Half 
of the carriers of the current and the most efficient ones have 
been used to form water molecules, and in the cell there remains 
only the slow-moving sodium ions and chlorine ions. If the 
acid and base in the respective layers were not quite equivalent 
in amount, a slight excess of one or the other will be shown by 
the litmus color, but the important part of the experiment, the 
serious loss of ions, will still be unmistakable from the great 
decrease in the conductivity of the solution between the plates. 
424. A Second Example of Class II: Action of HC1 on 
Sodium Acetate. — It will be recalled that acetic acid, HC 2 H 3 2 , 
neutralizes NaOH, forming sodium acetate, thus: 
HC 2 H 3 2 +NaOH^NaC 2 H 3 2 +H 2 0. 



Applications of the Ionic Hypothesis 259 

Acetic acid is a monobasic acid, only one of the four hydrogen 
atoms of each molecule being ionizable: 

HC 2 H 3 2 ^H++C 2 H 3 2 -. 

This acid is but little ionized in normal solution, the degree of 
ionization being only 0.4 per cent. On the other hand, solu- 
tions of its salts, like NaC 2 H 3 2 , are highly ionized: 

NaC 2 H 3 2 ^Na++C 2 H 3 2 -. 

If we mix equivalent amounts of HC1 and NaC 2 H 3 2 in 
solution we cannot see that any chemical change occurs; but 
that a reaction has occurred we may show convincingly with the 
help of the electrolytic cell, which is used to discover the change 
in conductivity during neutralization. In the lower layer this 
time we shall have sodium acetate and in the upper hydrochloric 
acid. As before, the lamp glows — both solutions are good con- 
ductors; the first by means of Na + and C 2 H 3 2 ~ ions, the 
second by means of H + and Cl~ ions. When we mix the two 
layers the decrease in brightness of the lamp shows that the 
conductivity has dropped off greatly, thus proving that many of 
the ions have been changed into non-conducting molecules. 
The compound equation is 

HC1±5H++C1- 

NaC 2 H 3 2 ^C 2 H 3 2 ~+Na+ 

It It 

HC 2 H 3 2 NaCl 

The graphs are shown in Figs. 60 and 61. Since of the four sub- 
stances concerned all but the acetic acid are highly ionized, while 
the latter is but little ionized, the reaction falls under Class II. 
When HO and NaC 2 H 3 2 solutions are mixed, the H + and 
C 2 H 3 2 ~ ions will unite far more completely than will any other 
pair of ions, and at the same time the molecules of HC1 and 
NaC 2 H 3 2 will continue to ionize until but very few remain 
(Fig. 61). Also Na + and Cl~ ions will unite partially to form 
molecules of NaCl. Therefore the equilibrium mixture will 
contain largely free acetic acid, for the most part un-ionized, 



260 



Introduction to General Chemistry 



together with common salt and its ions. Very little HC1 and 
NaC 2 H 3 2 will be present. 

425. Comparison of the First and Second Examples of Class 
II. — Fig. 60 shows the conditions of the solutions of hydrochloric 
acid and of sodium acetate before they are mixed, while Fig. 61 
shows the condition of the equilibrium mixture. These figures 
are almost a reproduction of Figs. 58 and 59, representing neutral- 
ization. In place of NaOH we have in the second case NaC 2 H 3 2 , 
which is also highly ionized; and in place of water we have 
HC 2 H 3 2 , which, like water, is but little ionized. In Fig. 61 the 
circle representing molecular acetic acid is nearly completely 



•x HCI 



. NaCH.O, 




QV 



\H + 



GNaC 2 H,0 2 ^""""xdHA-yunillUJiin, 




HC 2 HA ,-~-. 




NaCI 



Fig. 60 



Fig. 61 



cross-hatched, showing that the yield of this substance is nearly 
100 per cent. Although the reactions represented by Figs. 59 
and 61 are so nearly alike, there is a small difference due to the 
fact that acetic acid is ionized more than water. In consequence 
the formation of molecular acetic acid falls short of 100 per cent 
by a small fraction of 1 per cent. 

426. A Third Example of Class II: Action of NaOH on 
NH4CI. — Another important example of Class II is found in 
the action of sodium hydroxide and ammonium chloride. The 
addition of dilute NaOH to a solution of NH 4 C1 does not produce 
any visible effect; but evidence that the reaction 

NaOH+NH 4 Cl^NaCl+NH 4 OH 

takes place may be obtained in two ways: first, by finding a 
great decrease in conductivity on mixing superimposed layers 






Applications of the Ionic Hypothesis 261 

of the two solutions; and secondly, by noting the odor of 
ammonia given off by reason of the partial dissociation of the 
NH 4 OH present in the solution 

NH 4 OH±5NH 3 +H 2 0. 

The compound equation of the reaction follows: 

NaOH±5Na++OH- 

NH 4 C1^C1-+NH 4 + 

It It 

NaCl NH 4 OH. 

Since sodium hydroxide and ammonium chloride are highly 
ionized, and ammonium hydroxide is little ionized, this reaction 
is completely analogous to that between HC1 and NaC 2 H 3 2 : 

HCl+NaC 2 H 3 2 ^NaCl+HC 2 H 3 2 . 

Each reaction takes place nearly completely from left to right 
because one product is but little ionized. The graphs for 
this reaction, Figs. 62 and 63, are closely similar to those for 





Fig. 62 Fig. 63 

neutralization, Figs. 58 and 59, and for the reaction between 
HC1 and NaC 2 H 3 2 , Figs. 60 and 61. 

427. Summary of Class II Reactions. — As we have pointed 
out, all these reactions are alike, in that two highly ionized 
electrolytes react to form one highly ionized electrolyte and one 
little ionized electrolyte. Invariably reactions of this class are 
nearly complete. The smaller the degree of ionization of the 
little ionized product, the more completely the reaction takes. 



262 Introduction to General Chemistry 

In the resulting mixture the little ionized substance is present, 
of course almost wholly in the molecular form. 

428. Strength of Acids. — Since all salts are highly ionized, 
the reaction between any highly ionized acid and a salt of a 
little ionized acid must belong to Class II. We may therefore 
predict that, as in the second example studied, such reactions 
will give nearly 100 per cent yields of their products, and that 
in the resulting solution there will be present the little ionized 
acid instead of the highly ionized acid originally used. The 
highly ionized acid may be said to have displaced the little 
ionized acid from its salt. As a result, we may call the former 
a strong acid and the latter a weak acid, and may say that a 
strong acid always displaces a weak acid from its salts. 

429. Strength of Bases. — Just as we call a highly ionized 
acid a strong acid and a little ionized acid a weak acid, so we 
may call a highly ionized base a strong base and a little ionized 
base a weak base. Since all reactions between strong bases and 
the salts of weak bases (see third example, 426) are examples of 
Class II, we can predict that the yield of weak base and salt of 
the strong acid will be nearly 100 per cent. In other words, a 
strong base will always displace a weak base from its salt. 

430. Two Useful Laws. — The foregoing law and that given 
in the paragraph on the strength of acids (428) have been of 
very great practical convenience to chemists. These laws fail 
only when the salt of the weak acid is little ionized, a case so 
rare that the usefulness of the rules is virtually unaffected. The 
laws are of course only special cases of the fundamental one that 
if two highly ionized substances react to form one little ionized sub- 
stance and one highly ionized substance, the reaction will be nearly 
complete. 

431. Suppression of the Ionization of a Weak Acid or a 
Weak Base. — Since the strength of an acid or a base is deter- 
mined by its tendency to ionize, any factor that has an influence 
on this tendency will affect the strength or weakness of the acid 
or base. We must now consider this important subject and will 
begin by studying the action of NH 4 C1 on a solution of the weak 
base NH 4 OH. 



Applications of the Ionic Hypothesis 263 

If we add a little phenolphthalein to very dilute NH 4 OH a 
bright, red-colored solution results. This shows that the solu- 
tion is alkaline, and therefore that it contains an abundance of 
OH~ ions. Upon addition of a little NH 4 C1 to this red solution 
the color disappears almost completely. This proves that the 
number of OH~ ions present has been very greatly decreased. In 
order to understand how this has happened, we must consider 
the matter from the standpoint of ionic equilibrium. A solu- 
tion of NH 4 OH is ionized to a small extent, thus : 

NH 4 OH^NH 4 ++OH". 

Ammonium chloride, on the other hand, is very highly ionized: 
NH 4 C1^NH 4 ++C1". 

If then we add an equivalent amount of NH 4 C1 to a dilute solu- 
tion of NH 4 OH, the number of NH 4 + ions per cubic centimeter 
will be increased many fold. The OH~ ions present will therefore 
collide with NH 4 + ions and combine with them far more fre- 
quently than before. Since the rate of dissociation of NH 4 OH 
molecules into ions is not affected by the presence of the NH 4 C1, 
this increased rate of union of NH 4 + and OH~ ions causes a 
great shift to the left of the equilibrium 

NH 4 OH±5NH 4 ++OH-. 

For example, it has been found, by methods that we need not 
consider here, that the addition of 1 g. of NH 4 C1 to 100 c.c. of 

NIMH' '--sNH.+ ,<~- N 0H ' 






NHXI 
\ 

/ 




Fig. 64 Fig. 65 

decinormal NH 4 OH will decrease the number of OH~ ions 
present about one hundred fold. In other words, the ionization 
of the base will be decreased one hundred fold (see Figs. 64 and 
65). We may now state the general law of which the case just 



264 Introduction to General Chemistry 

studied is a typical example: The ionization of a weak base is 
greatly suppressed by the addition of a salt of the base. This 
means that a weak base is made still weaker by the addition of its 
soluble salts. 

In a similar manner the ionization of a weak acid is greatly 
suppressed by the addition of any of its soluble salts; that is, a weak 
acid is made still weaker by adding one of its salts. For example, 
the addition of NaC 2 H 3 2 to a red solution of acetic acid, 
HC 2 H 3 2 , containing litmus changes the color from red to 
purple, thus showing a great decrease in the number of H + ions, 
and therefore a great decrease in ionization of the acid. 

432. The Common Ion Law. — A base and any one of its 
salts must of necessity have one ion in common. (The Ntl 4 + 
ion is common to NH 4 OH and NH 4 CL) An acid also must have 
one ion in common with any of its salts. We may therefore 
state the principle of the foregoing laws as follows : Suppression 
of the ionization of a little ionized substance occurs when we add to 
its solution a highly ionized substance having a common ion. This 
is the Common Ion Law, a very important*generalization. The 
examples already cited are by no means the only ones of impor- 
tance. For example, it is plain that the ionization of NH 4 OH 
must be suppressed by the addition of NaOH or KOH because 
of the increase in concentration of the common OH~ ion; and 
that the ionization of HC 2 H 3 2 must likewise be suppressed by 
the addition of any strong acid like HO or HN0 3 . The effect 
of a highly ionized substance on the ionization of another highly 
ionized substance having one ion in common is of the same type 
but very much smaller in degree than when the second substance 
is slightly ionized. 

We shall next consider the application of the Common Ion 
Law to solutions of acids and bases and thus obtain a definition 
of the term neutrality. 

433. Neutrality Defined. — We have already learned (420) 
that water is slightly ionized, thus, 

H 2 O^H++OH-. 

Each cubic centimeter of pure water must therefore contain 
exactly as many H + as OH~ ions. Since all acids give H + 



Applications of the Ionic Hypothesis 265 

ions, the addition of an acid to water, in accord with the common 
ion law, will greatly suppress the ionization of water. Therefore 
acid solutions will contain far less OH~ ions per cubic centimeter 
than pure water. In an acid solution the number of H + ions 
greatly exceeds the number of OH~ ions. The ionization of 
water is also greatly suppressed by the addition of a base, since 
all bases have OH - ions in common with water. In basic solu- 
tions the number of H + ions per cubic centimeter is far less than 
in. pure water and therefore the number of OH - ions greatly 
exceeds the number of H + ions. Since we may consider water 
a typically neutral substance we may define a neutral solution as 
one in which the number of H + ions equals the number of OH~ 
ions. Since, as we have already learned, a strong acid completely 
neutralizes a strong base, as for example in the reaction 

HCl+NaOH^MaCl+H 2 0, 

we conclude that in the resulting solution the number of H + ions 
is just equal to the number of' OH~ ions: this is the criterion 
of complete neutrality. 

434. First Example of Class III: The Action of a Weak Acid 
on a Strong Base. — Under Class III we shall study reactions in 
which one little ionized and one highly ionized substance give 
products, one of which is little ionized, the other highly ionized. 
As the first example we shall study the reaction between little 
ionized acetic acid (a weak acid) and highly ionized sodium 
hydroxide (a strong base). These react thus: 

HC 2 H 3 2 +NaOH^H 2 0+NaC 2 H 3 2 . 

Of the products, water is very slightly ionized, while sodium 
acetate, NaC 2 H 3 2 , is highly ionized. If we mix equal volumes 
of normal solutions of HC 2 H 3 2 and NaOH, that is, if we add 
to the NaOH solution exactly that quantity of acetic acid that 
would neutralize it if the reaction were complete, we find that 
the resulting mixture is not neutral but is still alkaline to litmus. 
The fact that the mixture is alkaline means that the number of 
HO~ ions is greater than the number of H + ions present. The 



266 



Introduction to General Chemistry 



cause of this condition is most easily understood by aid of the 
compound equation 

HC 2 H 3 2 ^H++C 2 H 3 2 - 

NaOH±50H-+Na+ 

It it 

H 2 NaC 2 H 3 2 

and Figs. 66 and 67. At the instant of mixing, the solution 
contains an abundance of OH~ ions (Fig. 66); these reduce 
greatly the number of H + ions present by forming H 2 molecules: 

H++OH-^H 2 0. 

The removal of H + ions disturbs the equilibrium 
HC 2 H 3 2 ^H++C 2 H 3 2 -, 

which shifts greatly to the right, thus producing both H + and 
C 2 H 3 2 ~ ions. While the former unite with OH - almost (but 



HCH.0, 





Fig. 66 



it if 





c,H,or 



Fig. 67 



not quite) completely, the latter remain for the larger part free 
in the solution, and by their great tendency to unite again with 
H + ions to form little ionized HC 2 H 3 2 serve still further to 
diminish the number of free H + .ions. In the final equilibrium 
mixture, shown in Fig. 67, the number of OH~ ions is greater 
than the number of H + ions because of the great tendency of the 
latter to unite readily with either C 2 H 3 2 ~ ions or OH~ ions. 
That the OH~ ions get by far the lion's share of the H + ions is 
owing to the fact that water is much less ionized than acetic acid. 



Applications to the Ionic Hypothesis 267 

Since the mixture contains more OH~ than H + ions (see Fig. 
67) it is not neutral but alkaline. 

435. The Action of Water on Sodium Acetate. — In the fore- 
going paragraph we studied the equilibrium 

HC 2 H 3 2 + NaOH ±5 NaC 2 H 3 2 + H 2 . 

The composition of the equilibrium solution was shown in Fig. 67. 
It must be plain from the deduction of 281, that exactly the same 
equilibrium solution would be obtained if we should dissolve in 
the same quantity of water pure sodium acetate, NaC 2 H 3 2 , in 
exactly the amount that would be produced by the complete 
union of all the HC 2 H 3 2 and NaOH used in the first case. As a 
matter of fact we find that a solution of pure sodium acetate is not 
neutral but alkaline to litmus. The action of H 2 on NaC 2 H 3 2 
takes place thus: the salt first dissolves and at once ionizes 
highly to form many Na + and C 2 H 3 2 ~ ions. Water, although 
but slightly ionized, contains some H + and OH - ions. Occasional 
collisions of H + and C 2 H 3 2 ~ ions will occur, and part of these 
collisions will result in unions to form HC 2 H 3 2 molecules; and 
as the latter have but little tendency to ionize, the result is a 
great decrease in the number of H + ions present. This in turn 
disturbs the equilibrium 

H 2 O^H++OH-, 

which in consequence shifts to the right and so brings more OH~ 
ions into the solution. A few but not many of the OH~ ions 
unite with Na + ions to form molecules of NaOH, but most of 
the OH~ ions remain free, thus producing in the solution a 
decided excess of OH - ions over H + ions (see Fig. 67), and so 
making the solution alkaline to litmus. Briefly stated, water 
acts on sodium acetate to a small extent, thus, 

NaC 2 H 3 2 +H 2 O^HC 2 H 3 2 +NaOH, 

and since NaOH is highly ionized, while HC 2 H 3 2 is little ionized, 
the reaction of the solution is alkaline. The composition of a 
water solution of sodium acetate is that shown in Fig. 67. 

436. Hydrolysis of Salts. — The soluble salts of all weak acids 
with the strong bases sodium, potassium, calcium, or barium 



268 Introduction to General Chemistry 

hydroxide give alkaline solutions when dissolved in water. In 
every case the reason is the same as that given for the alkaline 
reaction of sodium acetate solution. The effect of water on the 
salt of a weak acid and a strong base is an example of the type 
of reaction called hydrolysis (or also hydrolytic dissociation). 
Hydrolysis may be defined as a double decomposition reaction 
in which water is one of the reacting substances. The solutions 
of salts of all weak acids and strong bases are alkaline in reaction. 
Other things being equal, the weaker the acid from which the salt 
is derived the greater the extent of the hydrolysis; that is, the 
greater the alkalinity of the solution. 

On the other hand, some salts (other than acid salts like 
NaHS0 4 ) give solutions that have an acid reaction (176). 
Among such are the chlorides, sulfates, and nitrates of copper, 
lead, iron, zinc, aluminum, etc. Experiments show that the 
hydroxides of all these elements are weak bases. It would 
therefore seem probable that the acidity of solutions of the salts 
of these bases with strong acids is due to hydrolysis, and that 
the behavior of such salts with water is the counterpart of the 
behavior of salts of weak acids with strong bases. 

437. A Second Example of Class III: The Action of a Strong 
Acid on a Weak Base. — The action of a strong acid on a weak 
base is plainly the reverse of that just discussed: the action of 
water on the salt of a strong acid and weak base. It follows that 
a weak base does not react completely with the theoretical or 
chemically equivalent amount of a strong acid, and in con- 
sequence the resulting mixture is still acid in its reaction. The 
action of HO on the weak base NH 4 OH will serve as a simple 

illustration : 

NH 4 OH^OH~+NH 4 + 
HC1^H++C1~ 

It It 
H 2 NH 4 C1. 

Comparison of this reaction with that for HC 2 H 3 2 and NaOH 
where we have a weak acid and strong base will bring out com- 
plete analogy. Experiment shows that a solution of NH 4 C1 in 
water is not neutral but slightly acid in reaction. Briefly 



Applications of the Ionic Hypothesis 



269 



stated, NH 4 OH does not completely neutralize an equivalent 
amount of HC1 because it is a weak base. Conversely, water 
acts on pure NH 4 C1 to form some free HC1 and NH 4 OH. 

438. Class IV: The Action of a Weak Acid and a Weak 
Base. — Under Class IV we shall include reactions between two 
little ionized substances, which give as products one little ionized 
and one highly ionized substance. The only reactions of im- 
portance in Class IV are those between a weak acid and a weak 
base, the products being water and a salt. Acetic acid and 
ammonium hydroxide are both moderately weak (but not 
extremely weak) . They react thus : 

HC 2 H 3 2 +NH 4 OH±5NH 4 C 2 H 3 2 +H 2 0. 

The reaction is not complete, as in the case of the action of a 
strong acid and a strong base, but reaches equilibrium when a 
few tenths of 1 per cent of the free un-ionized acid and free 
un-ionized base are still present in the solution. The conditions 
before and after the reactions are shown in Figs. 68 and 69. 





Fig. 68 



Fig. 69 



If on the other hand the solution is made by dissolving solid 
NH 4 C 2 H 3 2 in water partial hydrolysis takes place, giving a 
mixture the composition of which is also represented by Fig. 69. 

If both acid and base are extremely weak the extent of the 
hydrolysis will be much greater than in the case of NH 4 C 2 H 3 2 . 
In fact, in such cases hydrolysis may be so nearly complete that 
we may say that extremely weak acids in water solution do not 
form salts with extremely weak bases. 



270 Introduction to General Chemistry 

439. Heat of Ionization. — The heat liberated or absorbed by 
the complete dissociation into its ions of one formula of a dis- 
solved electrolyte is called its heat of ionization (cf. 366). In 
some cases heat is absorbed, in other cases it is liberated, when 
the substances are ionized, but in the great majority -of cases 
the heat of ionization is very small. For practical purposes we 
may say that the heat of ionization of readily ionizable elec- 
trolytes is almost negligible. Little ionized substances often 
have appreciable heats of ionization. This is notably the case 
with water, for which we have the following, 

H++OH~->H 2 0+ 13,700 cal. 

It was stated earlier (362) that the heat of neutralization of 

a strong acid by a strong base is almost the same in all cases, 

namely 13,700 calories. The reason can now be seen. We 

know that in the neutralization of a strong acid by a strong base 

in dilute solution the principal change is the union of H + and OH~ 

ions to form water. In other words, the simplified equation of 

neutralization is 

H++OH-->H 2 0. 

Since strong acids and bases, as well as most salts, have 
negligible heats of ionization; and since, moreover, very little 
dissociation or union of ions, other than H + and OH~, occurs 
in neutralization (421, Fig. 59), the heat produced in the reaction 
is simply that due to the formation of water from its ions. It 
is for this reason that heats of neutralization are practically 
the same for all strong acids and bases: 13,700 cal. for one 
formula weight (18 g.) of water formed. 

The heat of neutralization of ammonium hydroxide by a 
strong acid is 12,300 calories. The difference, 13,700—12,300 = 
1,400 cal., is the heat of ionization of the weak base. 

In reactions between solutions of two highly ionized salts 
which form by interaction two other highly ionized and easily 
soluble salts no appreciable heat change is observed. This is 
because in such reactions very little change takes place (418, 
Fig. 55), and such changes as do occur are accompanied by nearly 
negligible heats of ionization. 



Applications of the Ionic Hypothesis 271 

440. Indicators. — In addition to litmus, which is used so 
often to indicate the acidity or alkalinity of solutions, a number of 
other colored substances are also employed. These are called 
indicators. The more important indicators besides litmus are 
phenolphthalein and methyl orange. The former is a colorless 
substance which gives a bright red solution with alkalies. 
Methyl orange is orange color in neutral solution, pink in acid, 
and yellow in alkaline solution. In general, indicators are very 
complex chemical substances whose formulae need not be con- 
sidered at present. 

Since acid solutions always contain H + ions and alkaline 
solutions OH~ ions, we may say that an indicator is a substance 
which has one color in the presence of an excess of H + ions and 
a different color in the presence of an excess of OH~ ions. We 
might expect that every indicator would show its transition 
shade of color in an exactly neutral solution; that is, in a solu- 
tion where the number of H + ions equals the number of OH" 
ions. This, however, is not the case. In other words, most 
indicators do not indicate perfect neutrality. Litmus is a nearly 
perfect indicator, but phenolphthalein shows a change of color 
when the number of OH~ ions equals eighty times the number 
of H + ions; that is, if a solution contains more than eighty 
times as many OH~ as H + ions it colors phenolphthalein red 
(the alkaline color) ; if it contains less than eighty times as many 
OH~ ions as H + ions it leaves phenolphthalein colorless. On 
the other hand, methyl orange shows an orange color (its inter- 
mediate shade between pink, the acid color, and yellow, the 
alkaline color) when the number of H + ions is about a million 
times the number of OH - ions. Anomalous as it may seem at 
first thought, it is really fortunate that many of our indicators 
do not indicate perfect neutrality; for suppose we wish to dis- 
cover how much acetic acid a certain solution contains. We 
may titrate it accurately with normal or decinormal sodium 
hydroxide or other strong base if we use the right indicator 
(137) . Now we have learned that when acetic acid is mixed with 
exactly the theoretically equivalent amount of NaOH the result- 
ing solution is not perfectly neutral but in reality slightly alkaline 



272 



Introduction to General Chemistry 



(434). In accord with this we found that a solution of NaC 2 H 3 2 
was slightly alkaline to litmus, showing that the number of OH~ 
ions was greater than the number of H + ions. Therefore we 
must use as a titration indicator one which shows its change of 
color when the number of OH~ is greater than the number of H + 
ions. We find that phenolphthalein proves to be just right for 
the purpose. In general, we use phenolphthalein as indicator 
in titrating all moderately weak acids. 

If we wish to titrate NH 4 OH with HC1 we cannot use 
phenolphthalein, because a solution of NH 4 C1 contains more H + 
than OH~ ions. Such a solution " seems" acid to this indicator. 
We must use one which changes color when the number of H + 
ions exceeds the number of OH~ ions, and for this case we find 
methyl orange satisfactory. In general, we use methyl orange 
in titrating moderately weak bases. The acid used in such 
titrations must be a strong one. In titrating a strong base with 
a strong acid any of these indicators gives sufficiently accurate 
results. Table XVIII gives the colors of indicators in solutions 

TABLE XVIII 



Hydrogen ion concentration 
Hydroxyl ion concentration 

Methyl orange . 

Litmus 

Phenolphthalein 



IO~ 3 


io - s 


«,- 


10- 11 
Pink 
Red 
Colorless 


10-9 
Yellow 
Red 
Colorless 


10-7 
Yellow 
Purple 
Colorless 



IO" 8 

io- 6 
Yellow 
Blue 
Red 



of hydrogen and hydroxyl ion concentrations near those at 
which the color change occurs. In this table the concentrations 
are given in gram molecular weights per liter. If the H + con- 
centration is io~ 3 , 1,000 liters contain 1 g. of H + ion. 

441. Summary on Equilibrium between Soluble Electro- 
lytes. — If we mix solutions of two electrolytes, AB and CD, 
having no ion in common, a double decomposition reaction, 

AB+CD^AD+CB, 



takes place to a greater or less extent, because of the tendency 
of each positive ion to combine with each negative ion present. 



Applications of the Ionic Hypothesis 273 

If all four substances of the preceding equation are highly 
ionized (Class I, 415), the mixed solution will contain largely 
the four sorts of free ions, A + , B~, C + , and£>~. Only a small 
percentage of the dissolved material will be present as molecules. 
If all four substances are equally ionized, equal numbers of 
molecules of each will be present, as shown in Fig. 56.. 

If one of the four substances (say AD) is little ionized (Class 
II, 419), then the large numbers of A + and D~ ions shown in 
Fig. 56 cannot exist side by side in the mixed solution, since they 
will 'very largely combine to form AD molecules. The disap- 
pearance of A + and D~ ions allows A B and CD molecules more 
or less completely to dissociate. The final result, shown in 
Figs. 59, 61, and 63, is a nearly complete reaction, AD being 
present almost wholly in un-ionized form and CD to a small 
extent as molecules, but largely as C + and D~ ions. 

A generalization of much importance is found in the Common 
Ion Law: suppression of the ionization of a little ionize sub- 
stance occurs when we add to its solution a highly ionized 
substance having a common ion. 

Since in pure water the number of H + ions is equal to the 
number of OH - ions, and since we may consider pure water a 
perfectly neutral substance, we define a neutral solution as one 
in which the number of H + ions is exactly equal to the number 
of OH~ ions. 

Class III (434, 437) comprises reactions in which one little 
ionized substance reacts with a highly ionized substance to 
form products one of which is slightly, the other highly, ionized. 
Examples are found in the neutralization of a weak acid by a 
strong base; or of a weak base by a strong acid. In such cases 
the reaction is more or less incomplete. The weaker the acid 
or base taken, the less complete is the neutralization. Con- 
versely, salts of weak acids or of weak bases are hydrolyzed by 
water. The former give solutions which are alkaline, the latter 
those which are acid, in reaction. This kind of action is called 
hydrolytic dissociation. 

Under Class IV. (438) it was pointed out that weak acids and 
weak bases always react incompletely, and that when either 



274 Introduction to General Chemistry 

or both are extremely weak, salt formation may not occur in 
solution (177). 

We have seen that indicators change color according to the 
concentration of H + and OH~ ions present. Litmus shows 
its neutral tint when the numbers of H + and OH~ ions are 
nearly equal. Phenolphthalein requires an excess of OH~ 
ions to change color, while methyl orange requires an excess of 
H + ions. 



CHAPTER XIX 

APPLICATIONS OF THE IONIC HYPOTHESIS. REACTIONS 
INVOLVING CHANGES OF STATE 

442. Introduction. — In the present chapter we shall study 
precipitation from the standpoint of the ionic hypothesis in 
order to understand- the underlying principles of this most 
important means of separating substances. In equations for 
precipitation reactions, the substance precipitated will be indi- 
cated by a downward-pointing arrow. 

If we consider the familiar examples of precipitation repre- 
sented by the following equations, 

AgN0 3 +HCl±5AgCl|+HN0 3 , 
BaCl 2 +H 2 S0 4 ^BaS0 4 |+2HCl, 

we might conclude that AgCl and BaS0 4 are precipitated because 
they are insoluble in water. We might even be tempted to say 
that in the reaction 

AB+CD±^AD+CB, 

if either AD or CB is an insoluble substance it will be pre- 
cipitated. This statement contains something of the truth, 
but it is far from the whole truth, as the following examples will 
prove. Calcium carbonate, CaC0 3 (marble), is an almost 
insoluble substance. If we mix solutions of calcium chloride 
and carbonic acid we might expect to get a precipitate of calcium 
carbonate, thus, 

CaCl 2 +H 2 C0 3 ±> CaCO^-f- 2HCI. 

Not a trace of precipitate is formed. On the other hand potas- 
sium chlorate, KC10 3 , is easily soluble in water; but if we add 
a saturated solution of potassium bromide, KBr, to a saturated 
solution of sodium chlorate, NaC10 3 , a precipitate of KC10 3 
forms. Evidently the matter is not as simple as at first thought 
it appears to be. The separation of a solid from a solution is 

275 



276 Introduction to General Chemistry 

obviously the reverse of the passage of a solid into solution. 
Accordingly, in beginning the study of precipitation, it will be 
advisable for the student to read again sections 120-23. In 
section 122 it is stated, "A solution which at a fixed tempera- 
ture will dissolve no more of a given substance is called a 
saturated solution. When we speak of the solubility of a sub- 
stance we mean the amount of substance dissolved in a given 
amount of water in the case of the saturated solution." 

443. The Kinetic Theory of Solution. — When a soluble salt 
is brought into water its molecules begin to leave the surface 
of the solid and pass into the water. Immediately thereafter 
dissolved salt molecules will occasionally strike the surface 
of the solid and in some cases remain attached thereto. Finally, 
when the solution has become saturated we may imagine that 
the equilibrium between dissolved and solid salt is the result 
of the passage of molecules into and out of solution at exactly 

equal rates, thus: 

AB^AB 
Solid Dissolved 

This picture is, however, incomplete, since the salt is partly 
ionized. The dissolved molecules are therefore in equilibrium 
with their ions as well as with the solid salt, thus : 

AB^AB^A++B~ 
Solid Dissolved 

444. Graphic Representation of a Solid Electrolyte in 
Equilibrium with Its Saturated Solution. — We shall represent 

NaCl 

Fig. 70 

a solid electrolyte (acid, base, or salt) graphically by a cross- 
hatched square. The condition of a saturated solution of a 
soluble salt (NaCl, for example) in contact with an excess of 
the solid salt may then be represented as in Fig. 70. 





Applications of the Ionic Hypothesis 277 

445. The Solubility of Molecules. Molecular Solubility. — 

If we except a small number of electrolytes like sulfuric and 
nitric acids, which mix with water in all proportions, all other 
acids, bases, and salts have limited solubilities in water. Since 
all electrolytes are more or less ionized in solution, the dissolved 
substance is present partly as molecules and partly as ions. 
Therefore the total solubility of a substance in a solution saturated 
at a given temperature is the sum of the solubility of its molecules 
and the solubility of its ions. It seems reasonable to assume 
that the limited solubility of an electrolyte as a whole is the 
result of the limited solubility of its molecules rather than of its 
ions. Two reasons may be given for this assumption which will 
be amply confirmed by additional evidence to be considered 
later. 

In the first place the solid salt passes into and out of solution 
as molecules (see Fig. 70). If the molecules have a limited 
solubility, this would limit the solubility of the ions as well, since 
the latter and the former are directly in equilibrium. Therefore 
it is sufficient to assume limited solubility of the molecules in order 
to explain limited total solubility. Secondly, the small solubility 
of a difficultly soluble salt like CaS0 4 (100 c.c. of water dissolve 
o.25g. of CaS0 4 ) cannot be due to a correspondingly small 
solubility of Ca ++ or S0 4 ~~ ions, since solutions of CaCl 2 and 
many other easily soluble and highly ionized calcium salts con- 
tain an abundance of Ca ++ ions, and solutions of H 2 S0 4 and 
many easily soluble and highly ionized sulfates contain large 
concentrations of S0 4 ions. We shall assume, therefore, that 
at a given temperature the solubility of an acid, base, or salt is 
limited by the solubility of its molecules; and we shall call the 
solubility of the molecules (in the saturated solution) the 
molecular solubility (abbreviated M.S.) of the substance. Sum- 
marizing, we may say that when a solid electrolyte is mixed with 
water at a fixed temperature the substance dissolves and the 
concentration of the solution increases until the M.S. is reached; 
the solution is then saturated (at that temperature), and the 
molecules are in equilibrium with the ions and with the solid 
substance. 



278 



Introduction to General Chemistry 



- ~ v CD 
\ 




n 



w 



C8 



-v AO 



446. The Cause of Precipitation. — We are now ready to 
apply the foregoing principles, together with those learned in 
chapter xviii, to the process of precipitation. We have learned 
(414) that in the reaction 

AB+CD^AD+CB, 

if all four substances are easily soluble and highly ionized the 
resulting solution contains largely the four sorts of ions, together 

with small proportions of the 
four kinds of molecules. The 
conditions before and after 
the reaction are shown in Figs. 
55 and 56. Now let us suppose 
that one of the products AD 
is not very soluble, so that its 
molecular solubility (M.S.) is 
less than that corresponding to 
the concentration of the mole- 
cules of AD formed in reaction 
(Fig. 56). In this case the 
amount of molecules of AD 
formed in excess of the M.S. 
will separate out of solution as 
a precipitate. Fig. 71 shows 
the resulting condition for the case where the M.S. of AD is 
rather small but not extremely small. By comparison of Figs. 
56 and 71 we see that an appreciable shift in equilibrium of the 
dissolved substances accompanies the partial precipitation of AD. 

447. The Precipitation of KC10 3 . — An actual example con- 
forming perfectly to the conditions set forth in the preceding 
paragraph is found in the reaction 

KBr+NaC10 3 ^KC10 3 |+NaBr. 

Of the four salts, all are very soluble except KC10 3 , which dis- 
solves only to the extent of 7 g. in 100 c.c. of water at 18 . All 
four salts are highly and about equally ionized in solutions of 
equal concentration. The conditions of the solutions of KBr 



u 




Fig. 71 



Applications of the Ionic Hypothesis 279 

and NaC10 3 before mixing are shown with sufficient accuracy 
by Fig. 55, while Fig. 56 shows the condition which the mixed 
solution would reach if KC10 3 were also very soluble. It hap- 
pens, however, that the amount of molecular KC10 3 which 
tends to be formed exceeds the rather small M.S. of this sub- 
stance, and in consequence the excess above the M.S. separates 
as a precipitate. Precipitation continues until the amount of 
molecular KC10 3 left in solution is equal to the M.S. of this sub- 
stance. The mixture is then in the condition of equilibrium 
shown in Fig. 71. Comparison of Figs. 56 and 71 shows that 
the removal (by precipitation) of KC10 3 from the solution causes 
a marked shift in the equilibrium. We may trace the stages as 
follows: Fig. 56 shows the condition that would exist if no 
precipitation occurred. The removal of KC10 3 , results in the 
further union of K + and C10 3 ~~ ions to form more KC10 3 . The 
resulting loss of K + and C10 3 ~ ions promotes the further ioniza- 
tion of KBr and NaC10 3 respectively and thus increases the 
numbers of Br" and Na + ions. The latter ions unite in part 
to form additional molecular NaBr. The final result is the 
change from the condition of Fig. 56 to that of Fig. 71. The 
principles here exemplified apply to all double decomposition 
precipitations. 

448. The Precipitation of CaC0 3 . — Let us now consider a 
case in which one of the products is precipitated almost com- 
pletely. The reaction 

CaCl 2 +Na 2 C0 3 ^CaC0 3 |+2NaCl, 

in which CaC0 3 is the precipitate, will serve as a typical illustra- 
tion. Although CaC0 3 appears to be insoluble in water, it is in 
fact slightly soluble, and has therefore a definite but very small 
M.S. The other three salts, CaCl 2 , Na 2 C0 3 , and NaCl, are 
easily soluble and highly ionized, and in consequence the reaction 
between solutions of CaCl 2 and Na 2 C0 3 tends to reach the con- 
dition shown in Fig. 56 illustrating a Class I reaction (414). In 
this respect it completely resembles the reaction 

KBr+NaC10 3 ^KC10 3 ^+NaBr. 



28o 



Introduction to General Chemistry 



' ~ - Na 2 C0, 



CaCI 2 




co»- 



NCa^ 



It differs however from this reaction in that the M.S. of CaC0 3 
is extremely small compared with the M.S. of KC10 3 . In con- 
sequence the CaCQ 3 formed pre- 
cipitates almost completely, as 
illustrated in Fig. 72. In all 
double decomposition reactions 
of the above-mentioned types 
(all involved substances highly 
ionized) the precipitation is the 
more complete the less the M.S. 
of the precipitate. 

449. The Action of H 2 C0 3 
on CaCl 2 . — In section 442 it 
was pointed out that H 2 C0 3 
does not precipitate CaCl 2 , as 
we might expect according to 
the following hypothetical 
equation: 



w 



^.NaCI 



H 



H 



CaCO, 




CaC0 3 



Fig. 72 



CaCl 2 +H 2 C0 3 ^CaC0 3 -|+2HCl. 



The reason is as follows: carbonic acid H 2 C0 3 is a very weak 
acid and in consequence yields but very few C0 3 ions; and 
although CaCl 2 gives an abundance of Ca ++ ions, the concen- 
tration of C0 3 ions is so small that the concentration of 
CaC0 3 molecules formed is less than the M.S. of this substance. 
Therefore no precipitation of CaC0 3 takes place. The differ- 
ence in behavior of H 2 C0 3 and Na 2 C0 3 toward a solution of 
CaCl 2 is wholly due to the difference in their tendencies to ionize, 
in consequence of which a solution of H 2 C0 3 contains exceedingly 
few C0 3 ~~ ions as compared with a solution of Na 2 C0 3 . 

The behavior of H 2 C0 3 is typical of that of all weak (little 
ionized) electrolytes. In the precipitation of salts, weak acids 
and bases are, in general, less efficient precipitants than their 
salts, since the latter are highly ionized. 

450. The Precipitation of Magnesium Hydroxide, Mg(OH) 2 . 
— We shall next discuss in detail the precipitation of Mg(OH) 2 , 



Applications of the Ionic Hypothesis 281 

not so much because of the chemical importance of this sub- 
stance, but because the reactions illustrate in a striking way some 
of the most important principles of ionic equilibrium. 

If we add NaOH to a solution of magnesium chloride, 
MgCl 2 , we obtain an abundant white precipitate of Mg(OH) 2 , 
formed as follows: 

MgCl 2 +2NaOH^Mg(OH) 2 ^+ 2 NaCL 

If we use NH4OH instead of NaOH the reaction is similar but 
reaches a state of equilibrium when only a part of the magnesium 
is precipitated : 

MgCl 2 + 2 NH 4 OH ±> Mg(OH) a >H- 2 NH 4 C1. 

If we add to a MgCl 2 solution a solution of NH 4 OH mixed with 
sufficient NH 4 C1, no precipitation occurs. We shall now explain 
these facts. In the first place we may say that the action of 
NaOH on MgCl 2 is analogous to the action of Na 2 C0 3 on CaCl 2 , 
the two equations being 

MgCl 2 + 2 NaOH ±^ Mg(OH) 2 ^+ 2 NaCl, 
CaCl 2 +Na 2 C0 3 ^CaC0 3 |+2NaCl. 

Both MgCl 2 and NaOH, like CaCl 2 and Na 2 C0 3 , are easily solu- 
ble and highly ionized; sodium chloride, one of the products in 
both reactions, is also highly ionized. The other product in the 
first reaction, Mg(0H) 2 , is but slightly soluble, and like CaC0 3 
it is therefore precipitated almost completely. 

If, however, we use NH 4 OH instead of NaOH, the reaction is 
far from complete. The reason can best be seen by the aid of 
Figs. 73 and 74. Figure 73 shows the condition of the solutions 
before they are mixed; Fig. 74 represents the condition of the 
mixture. 

It will be recalled, as shown in Fig. 73, that NH 4 OH is but 
little ionized. Still its solution yields sufficient OH~ ions to 
produce in reaction with magnesium chloride solution more 
Mg(0H) 2 than the M.S. of the latter difficultly soluble substance. 
The excess of Mg(0H) 2 above its M.S. precipitates, Fig. 74. As 
these changes go on, molecules of NH 4 OH continue to ionize, thus 



282 



Introduction to General Chemistry 



bringing into the solution far more NH 4 + ions than were 
originally present (cf. Figs. 73 and 74). The presence of the 
large excess of NH 4 + ions restricts the number of OH~ ions to 
such an extent that a state of equilibrium is reached in reaction, 

Mg(OH) 2 (dissolved)±5Mg+++ 2 OH-, 

while there is still a considerable amount of magnesium in the 
form of Mg ++ ions and MgCl 2 molecules left in the solution. 
After this condition is reached 
no more Mg(OH) 2 precipi- 
tates. We therefore conclude 
that NH 4 OH precipitates 
Mg(0H) 2 only partially, (1) 
because NH 4 OH is a weak or 
little ionized base, and (2) 



NlgCl, 





NH« + 



0H- 



t- ' 



u 



H 





NH«0H 



T 1 



,NH 4 + ,--- s 0H- 



1 ^ | 


Ig(OH), 


/ \ / 




W 




1 i 
1 l 


Mg(0H) s 


1 ! 
■ I 





Fig. 73 



Fig. 74 



because the accumulation of NH 4 + ions (Fig. 74) finally restricts 
the OH~ concentration to so small a value that the Mg ++ 
and OH~ ions are just in equilibrium with the amount of 
Mg(0H) 2 corresponding to its M.S. It will now be easy to 
understand why no Mg(0H) 2 is precipitated when NH 4 OH, 
mixed with considerable NH 4 C1, is added to a MgCl 2 solution. 

The NH 4 C1 furnishes at once such an excess of NH 4 + ions 
that the OH~ concentration is decreased to so small a value that 
less Mg(0H) 2 is formed in the reaction 

Mg+++20H-±5Mg(OH) 2 

than corresponds to its M.S. Therefore no precipitation occurs. 

451. The Precipitation of Ferric Hydroxide, Fe(OH) 3 . — The 

reaction 

FeCl 3 -r-3NH 4 OH±>Fe(OH) 3 NH-NH 4 Cl 



Applications of the Ionic Hypothesis 283 

takes place with the practically complete precipitation of brown 
Fe(OH) 3 , which is almost insoluble in water. The completeness 
of precipitation is not noticeably affected by the addition of much 
NH 4 C1. There are two reasons for the difference in behavior of 
Fe(OH) 3 and Mg(0H) 2 : (1) the former is much more insoluble 
in water than the latter, so that the M.S. of the latter (although 
small) is perhaps 1,000 times as large as that of the former; 
(2) Mg(0H) 2 is a rather strong (highly ionized) base, while 
Fe(OH) 3 is a very weak base. Therefore in the reactions 

Mg+++20H-±5Mg(0H) 2 (dissolved), 
Fe++++30H-± 5 Fe(OH) 3 (dissolved), 

for equal concentrations of Mg ++ , Fe +++ , and OH~ ions far 
less Mg(OH) 2 is formed than Fe(OH) 3 . The presence of NH 4 C1 
decreases the OH - concentration of NH 4 OH, but not sufficiently 
to prevent the practically complete precipitation of Fe(OH) 3 
because of the weakness of the latter and its exceedingly small 
M.S. The weaker a base and the smaller its M.S., the more com- 
pletely is it precipitated by N HfiH, and the less its precipitation 
is hindered by the presence of ammonium salts. 

452. Classification of Precipitations. — The various examples 
of precipitation just studied cover the important principles con- 
cerned. We may now cite a few additional examples of each of 
these classes of precipitation. In the reaction 

KBr+NaC10 3 ^KC10 3 ^+NaBr, (447) 

all four substances are highly ionized, and all but KC10 3 are very 
soluble. The latter is partially precipitated because it is formed 
in excess of its not very large M.S. The following reactions are 
of this class: 

CaCl 2 +2NaC10 3 ^2NaCl|+Ca(C10 3 ) 2 , 
Pb(N0 3 ) 2 +2NaCl^PbCU+2NaN0 3 , (167) 

CaCl 2 +Na 2 S0 4 ^CaS0 4 ^+2NaCL (153) 

In the first reaction saturated solutions are required to give a 
precipitate of NaCl. 



284 Introduction to General Chemistry 

In the second example studied, 

CaCl 2 +Na 2 C0 3 ^CaC0 3 |+2NaCl, (448) 

the precipitate CaC0 3 has an extremely small M.S. Its pre- 
cipitation by Na 2 C0 3 is almost complete. Other reactions of 
this class are: 

AgN0 3 +NaCl = AgCty+NaN0 3> (382) 

AgN0 3 +KBr = AgBr|+KN0 3 , (257) 

Pb (N0 3 ) 2 + CuS0 4 = PbS0 4 ^+ Cu(N0 3 ) 2 , (167) 

BaCl 2 +Na 2 S0 4 = BaS0 4 ^+ 2NaCl, (164) 

MgCl 2 + 2 NaOH = Mg(OH) 2 |+ 2 NaCl. (450) 

In the third example (449) it was shown that H 2 C0 3 did not 
give with CaCl 2 a precipitate of CaC0 3 , because the former is 
very little ionized and therefore yields very few C0 3 ~ ~ ions. The 
following pairs of substances also fail to give precipitates, because 
in each case of the weakness of the acid coupled with the moderate 
solubility of the salt, that might be precipitated: 

CaCl 2 and H 3 P0 4 , 
FeCl 2 and H 2 S, 
AgN0 3 and HC 2 H 3 2 . 

On the other hand the reactions 

3 CaCl 2 +2Na 3 P0 4 = Ca 3 (P0 4 ) 2 |+6NaCl, (158) 

FeCl 2 + (NH 4 ) 2 S = FeS^+ 2NH 4 C1, 
AgN0 3 +NaC 2 H 3 2 = AgC 2 H 3 2 |+NaN0 3 , 

give abundant precipitates, because, instead of the weak acids, 
we use their salts, which are highly ionized. 

The fourth example, which belongs to Class II, was taken up 
in contrast to the fifth example, which is typical of a fourth class 
of precipitation reactions. The fourth and fifth examples were: 

MgCl 2 +2NaOH±*Mg(OH) 2 >H-2NaCl, (450) 

MgCl 2 +2NH 4 OH±>Mg(OH) 2 ^+2NH 4 Cl. (450) 

The precipitation of Mg(0H) 2 is nearly complete in the first 
reaction but only partial in the second, owing to the moderate 
solubility of Mg(OH) 2 and the little ionization of NH 4 OH, 
especially in the presence of its salts. 



Applications of the Ionic Hypothesis 285 

The following reaction of a weak electrolyte (H 2 S) also results 
in partial precipitation: 

ZnCl 2 +H 2 S^ZnS|+2HCl. 

In this case the precipitation is prevented by the presence of 
an excess of HC1 or other strong acid, because of the suppression 
of the ionization of the H 2 S by the H + ions of the strong acid. 
The sixth example dealt with the precipitation of Fe(OH) 3 : 

FeCl3+3NH 4 OH± 5 Fe(OH) 3 |+3NH 4 Cl. (451) 

In this case the precipitate is so insoluble (M.S. so small) and 
so weak (little ionized) that it is practically completely pre- 
cipitated by NH 4 OH even in the presence of NH 4 C1. Although 
we call NH4OH a weak base, it is enormously stronger than 
Fe(OH) 3 , even when mixed with much NH 4 C1. Other reactions 
which fall into this class are : 

A1C1 3 +3NH 4 0H^A1(0HU+ 3 NH 4 C1, (174) 

CrC] 3 + 3 NH 4 OH^Cr(OH) 3N H- 3 NH 4 Cl, (344) 

CuS0 4 +H 2 S^CuSnH-H 2 S0 4 , 

Ag 2 S0 4 +H 2 S^Ag 2 S|+H 2 S0 4 . 

Excess of NH 4 C1 in the first two cases, and of HC1 or H 2 S0 4 in 
the last two cases, fails to prevent practically complete precipita- 
tion. 

453. Precipitation by Adding a Substance Having a Common 
Ion. — We have learned in the foregoing chapter (432) that if we 
add to the solution of an electrolyte, AB, enough of another 
highly ionized electrolyte, AC, having a common ion, A, to 
increase the concentration of the common ion, the degree of 
ionization of the first substance will be suppressed. If now the 
substance A B is not very soluble, the suppression of its ioniza- 
tion caused by adding AC may increase the concentration of 
the A B molecules to such an extent as to exceed the M.S. of AB. 
In consequence part of AB will separate out as a precipitate. 
For example, if a few bubbles of HC1 gas are passed into a 
saturated solution of NaCl, a precipitate of NaCl is formed. A 
similar result is also produced by adding a little concentrated 
HC1 to a saturated salt solution. In each case the ionization 



286 Introduction to General Chemistry 

of the salt is suppressed by reason of the increase in concentra- 
tion of the Cl~ ions, and the concentration of the molecular NaCl 
increased beyond the M.S. of this substance. Salt precipi- 
tates until the concentration of molecular NaQ falls to the 
value corresponding to its M.S. Another example illustrating 
the same principle is found in the precipitation of KC10 3 from 
its saturated solution by the addition of a saturated solution 
of either KBr or NaC10 3 . We may say that as a general rule the 
lotal solubility of a salt {molecular and ionic) is diminished by the 
presence in the solution of another electrolyte having a common ion. 
454. Conditions Favoring Precipitation. — In the reaction 

AB+CD=AD+CB 

precipitation will occur if one of the products, say AD, is formed 
as molecules in greater concentration than its molecular solu- 
bility. In brief, if the M.S. is exceeded, precipitation will occur. 
Now the smaller the M.S. of AD, the more probably will it be 
exceeded. 

On the other hand the M.S. is the more likely to be exceeded 
the greater the concentration of AD which tends to be produced 
in the reaction. The various factors which determine the 
amount of AD produced (when AD is soluble) have been dis- 
cussed at length in chapter xviii (Summary, 441). 

These applications of the ionic hypothesis have the following 
bearing on the practice of precipitation. In the first place, if 
we are to precipitate from solution an "insoluble" salt of a weak 
acid we use as the precipitant a soluble salt of the weak acid 
instead of the acid itself, since the former is highly ionized, 
while the latter is not. (The term precipitant means the reagent 
added to cause precipitation.) If, however, we are to precipitate 
an insoluble chloride we may use either a soluble chloride or 
hydrochloric acid, since this strong acid is as highly ionized as 
its salt. When the precipitant is added to a given solution, a 
precipitate may not appear until considerable reagent has been 
added. When it is no longer possible to see if more precipitate 
is forming with further additions of the reagent, a small portion 
of the mixture is filtered, or, better, the precipitate is allowed 



Applications of the Ionic Hypothesis 287 

to settle, and the clear solution is tested with more of the pre- 
cipitant. Only moderate excesses of the precipitant are used 
as a rule, since in many cases the precipitant reacts farther with 
the precipitate to form new and soluble compounds, with the 
result that the precipitate dissolves in an excess of the reagent 
added. 

455. Dissolving Solid Substances. — Substances which are 
not readily soluble in water often dissolve easily in solutions of 
other electrolytes. In such cases we may imagine that chemical 
reaction gives rise to new products which are soluble in water. 
Here is a case in point: Calcium hydroxide, Ca(0H) 2 , is but 
slightly soluble in water (o.i2g. in iooc.c), giving a very 
dilute solution known as limewater. If we mix a few grams of 
Ca(OH) 2 with 100 c.c. of water, most of the solid remains undis- 
solved. If now we add dilute HC1 to the mixture, the solid 
finally passes completely into solution. The explanation is as 
follows: The small amount of dissolved Ca(0H) 2 (which is a 
strong base) is neutralized by the added HC1 to form very soluble 

Ca(OH) 2 + 2HC1^ CaCl 2 + 2 H 2 0. 

More Ca(0H) 2 then dissolves in the water in the tendency to 
keep the concentration of the dissolved Ca(0H) 2 up to its M.S.: 

Ca(OH) 2 ±> Ca(OH) 2 ±5 Ca+++ 2OH- 
Solid Dissolved 

As fast as Ca(0H) 2 passes into solution it reacts with the HC1 
present. If the chemically equivalent amount of HC1 is added, 
all Ca(OH) 2 will finally dissolve, and the solution will consist 
simply of CaCl 2 dissolved in water. 

A perfectly analogous reaction is found in the dissolving of 
the difficultly soluble, strong base Mg(OH) 2 in dilute HC1: 

Mg(0H) 2 + 2HCI ±> MgCl 2 + 2 H 2 0. 

Even if the base is weak and much less soluble than either 
Ca(0H) 2 or Mg(0H) 2 , it will usually dissolve in water upon the 
addition of a strong acid. For example, Fe(OH) 3 is a very weak 



288 Introduction to General Chemistry 

base almost insoluble in water; it dissolves readily in dilute 
HO, forming a solution of ferric chloride, 

Fe(OH) 3 + 3 HCl^FeCl 3 + 3 H 2 0. 

The stages in the process of dissolving may be considered to be 
comparable to those in the case of the dissolving of Ca(0H) 2 in 
dilute HCL 

Most bases, with the exception of the hydroxides of sodium, 
potassium, ammonium, and barium, are very little soluble 
in water. All such so-called insoluble bases dissolve in dilute 
HC1, HN0 3 , an d H 2 S0 4 to form clear solutions, if their cor- 
responding salts with these acids are soluble in water. 

456. Dissolving Little Soluble Salts of Weak Acids by Solu- 
tions of Strong Acids. — Silver acetate is a rather difficultly solu- 
ble salt (i.og. dissolves in 100 c.c. H 2 at 18 ) which is easily 
made by precipitating AgN0 3 with NaC 2 H 3 2 , 

AgN0 3 +NaC 2 H 3 2 ^AgC 2 H 3 2 +NaN0 3 . (452) 

If we mix 3 or 4 g. of AgC 2 H 3 2 with 100 c.c. of H 2 0, only a small 
portion dissolves; but upon addition of dilute HN0 3 the whole 
of the solid passes into solution. Silver acetate is the salt of 
the weak acid HC 2 H 3 2 , and, as we have already learned (428), 
a strong acid reacts more or less completely with the (soluble) 
salt of a weak acid to form the weak acid and the salt of the strong 
acid. This was shown earlier in the case of the reaction 

HCl+NaC 2 H 3 2 ^HC 2 H 3 2 +NaCl. (424) 

Nitric acid reacts similarly with the dissolved portion of the 
AgC 2 H 3 2 , 

HN0 3 +AgC 2 H 3 2 ^HC 2 H 3 2 +AgN0 3 . 

The reaction is nearly complete, and both products are easily 
soluble. The dissolved molecular AgC 2 H 3 2 being thus removed 
from the solution, more of the solid passes into solution in the 
tendency to keep the concentration of molecular AgC 2 H 3 2 up 
to its M.S. But as this salt reacts with the HN0 3 present as 



Applications of the Ionic Hypothesis 289 

soon as it comes into solution, its M.S. is never reached, so that 
finally all of the solid passes into solution. The solution con- 
sists largely of AgN0 3 and its ions, together with molecular 
acetic acid. 

In many other cases so-called "insoluble" salts of weak acids 
dissolve in solutions of strong acids like HC1, HN0 3 , and H 2 S0 4 . 
The following reactions are of this type: 

FeS+2HCl = H 2 S-f-FeCl 2 , 

CaC0 3 + 2HCI = H 2 C0 3 + CaCl 2 , 

Ca 3 (P0 4 ) 2 +6HCl = 2H 3 P0 4 + 3 CaCl 2 . 

However, not all " insoluble" salts of weak acids dissolve in 
strong acids. For example, CuS, which comes down as a black 
precipitate when H 2 S is passed into a solution of a copper salt, 
and is therefore a salt of the very weak acid H 2 S, does not dissolve 
appreciably in cold HC1. The reason for this is directly trace- 
able to the extremely small M.S. of CuS. In general the smaller 
the M.S. of a salt of a weak acid the less soluble it is in a strong 
acid. Other examples of this sort are found in Ag 2 S and HgS, 
neither of which is dissolved appreciably by dilute HC1 or 
H 2 S0 4 . 

457. Weak Acids and Salts of Strong Acids. — We have 
already learned (282) that the equilibrium mixture has the same 
composition whether we start with one pair of substances of a 
reaction or the equivalent amounts of the other pair. In accord 
with this principle we always find that if a reaction takes place 
practically completely in one direction, the reverse of the reaction 
does not succeed under the same conditions of temperature and 
concentration. In sections 449 and 452 it was stated that 
mixtures of the following pairs of substances fail to give pre- 
cipitates, although little soluble salts would be formed by 
double decomposition : 

H 2 C0 3 and CaCl 2 , 
H 3 P0 4 and CaCl 2 , 
H 2 S and FeCl 2 , 
HC 2 H 3 2 and AgN0 3 . 



290 Introduction to General Chemistry 

Therefore we may be certain that calcium carbonate, calcium 
phosphate, and ferrous sulphide are soluble in hydrochloric acid, 
and that silver acetate is soluble in nitric acid. Since the deter- 
mining factor in dissolving each of these salts is the formation 
of the weak acid, we may go farther and predict that any strong 
acid will dissolve these salts. Sometimes a new insoluble salt 
is formed by the strong acid, as when hydrochloric acid acts on 
silver acetate; but such reactions are secondary to the solution 
of the original salts. 

458. "Insoluble" Salts of Strong Acids.— The "insoluble" 
salts of strong acids are not as a rule dissolved to an appre- 
ciable extent by solutions of other strong acids. For example, 
AgCl is not appreciably dissolved by HN0 3 , although the 
products HC1 and AgN0 3 of the hypothetical reaction 

AgCl+HN0 3 ^HCH-AgN0 3 

are both easily soluble substances. A reaction in which both 
of the products are highly ionized, as in this case, falls in Class I 
(414). In all such reactions very little chemical change occurs, 
and this is more strikingly true the more dilute the solution. As 
we are now considering the case where one of the substances 
taken is nearly insoluble in water, the solution of this substance 
must be exceedingly dilute. Comparing the action of HN0 3 
on AgC 2 H 3 2 and AgCl, we may say that the first reaction takes 
place readily because of the tendency of H + and C 2 H 3 2 - to 
unite nearly completely to form little ionized HC 2 H 3 2 ; and 
that the second reaction does not progress far because of the 
slight tendency for H + and Cl~ ions to unite, since HC1 is 
nearly completely ionized in very dilute solution. 

459. Evolution of a Gas. — Substances may separate from 
solutions in two ways: (1) as solid precipitates and (2) as gases. 
We have considered the first case and shall now take up the 
second, and we shall see that if a substance separates from a 
solution as a gas the effect on the ionic equilibrium is the same 
as if the substance separated as a solid. The principles that 
apply to precipitation apply also, with slight obvious modifica- 
tions, to gas evolution. Gases have limited solubilities, and 



Applications of the Ionic Hypothesis 291 

instead of the M.S. of the precipitate we have the molecular 
solubility (M.S.) of the gas. Let us now consider a few well- 
known reactions as illustrations. 

460. The Action of H 2 S0 4 on NaCl. — If we mix dilute solu- 
tions of H 2 S0 4 and NaCl no visible effect is produced, although 
in the solution the reaction 

H 2 S0 4 -f-NaCl^HCl+NaHS0 4 

takes place partially. This is a Class I reaction (414) since all 
four substances are easily soluble and highly ionized. Therefore 
the dilute solution contains chiefly the ions and relatively few 
molecules. Nevertheless some HC1 molecules are formed even 
in dilute solution, but, as HC1 is a very soluble gas, none of it 
escapes from the solution. 

On the other hand the results are quite different if but little 
water is present. In making hydrochloric acid (103) 58 g. of 
NaCl, 100 g. of concentrated H 2 S0 4 , and 30 g. of water were 
mixed in a flask and gently heated. 

The proportions of NaCl and H 2 S0 4 taken were those indi- 
cated by the foregoing equation, since the 100 g. of concentrated 
acid taken consists of 98 g. of actual H 2 S0 4 and 2 g. of H 2 0. 
If the reaction should take place completely, 36. 5 g. of HC1 
would be formed. This is far more HC1 than the 32 g. of water 
present can hold in solution, especially when the mixture is 
heated. Therefore HC1 gas escapes from the solution. The loss 
of HC1 by the solution impedes the reverse action on the 
NaHS0 4 present and so causes a great shift to the right of the 
equilibrium that would otherwise be reached in the reaction 

H 2 S0 4 +NaCl^HClf+NaHS0 4 . 

As a consequence this reaction goes nearly completely from left 
to right under the conditions described (103), the HC1 being 
given off as gas. In equations for reactions involving gas 
evolution the gas will be indicated by an upward-pointing arrow. 

461. The Action of HC1 on CaC0 3 . — We have already seen 
that carbonic acid, H 2 C0 3 , does not precipitate CaC0 3 from a 
CaCl 2 solution, and have learned that this is because H 2 C0 3 is 



292 Introduction to General Chemistry 

so little ionized that insufficient molecular CaC0 3 is formed to 
exceed its M.S. This fact indicates that the reactions 

2HCI+ CaC0 3 ±» CaCl 2 + H 2 C0 3 , 
H 2 C0 3 ^H 2 0+C(VK 

will take place practically completely, since in all reactions 
between electrolytes exactly the same proportions of the same 
products result, whether we start with one pair of substances 
or the chemically equivalent amounts of the other pair (282). 
The dissolving of CaC0 3 in dilute HC1 takes place as follows: 
CaC0 3 first dissolves to the limit of its (very small) M.S. in the 
water present ; the dissolved molecules then ionize rather highly : 

CaC0 3 ±?Ca+++C0 3 ~. 

The C0 3 ions unite nearly completely with the H + ions of 
the highly ionized HC1 present, 

2 H++C0 3 --^H 2 C0 3 , 

and to a small extent Ca ++ and Cl~ ions unite to form (easily 
ionized) CaCl 2 . The nearly complete removal of C0 3 ~~ ions 
allows the further ionization of CaC0 3 , and this change permits 
the passage of more solid CaC0 3 into solution. The quantitative 
relations are such that these changes continue until all CaC0 3 
has dissolved. Incidentally the H 2 C0 3 , which is unstable, dis- 
sociates, to a large extent, into water and C0 2 , 

H 2 C0 3 ^H 2 0+C0 2 f, 

and, as the latter is not very soluble, much of it passes off as 
a gas. 

The several reactions are shown in the following diagram : 

C0 2 (gas) 
It 
H 2 C0 3 ^H 2 0+C0 2 

It 

2HC1^2C1- + 2H+ 

CaC0 3 ±5CaC0 3 ^Ca+++C0 3 ~ 

(Solid) If 

CaCL 



Applications of the Ionic Hypothesis 293 

462. The Action of NaOH on NH 4 C1. — Another example of 
gas evolution, which, however, does not introduce any new 
principle, is found in the reaction 

NaOH+NH 4 Cl = NaCl+NH 4 OH, 

which takes place more or less completely when solutions of the 
two initial substances are mixed. This reaction was studied 
under Class II (426), where it was pointed out that it takes place 
nearly completely because NH 4 OH is a weak base. This base 
is also unstable, readily dissociating, thus, 

NH 4 OH^H 2 0+NH 3 t, 

and, since NH 3 is a gas, it will in part escape from the solution. 
The more concentrated the solution and the higher the tempera- 
ture the more completely will the NH 3 be evolved as gas. The 
loss of NH 3 from the solution promotes the dissociation of 
NH 4 OH, and this in turn favors a further shift in equilibrium 
from left to right in the main reaction. The various reactions 
and their relations are fully shown in the following diagram : 

NH 3 (gas) 
It 
NH 4 OH±^H 2 0+NH 3 
It 
NH 4 C1^C1"+NH 4 + 

NaOH^Na++OH~ 

it 
NaCl 

463. The Factors Governing Gas Evolution. — The various 
factors which are favorable to ,gas evolution are very similar 
to those which were found to favor precipitation, although there 
are some differences aside from the fact that in the one case we 
are dealing with a gas and in the other with a solid product. If 
one of the products of a reaction is gaseous it will be given off 
from the solution, the more completely, the larger the propor- 
tion of it formed in the reaction and the less soluble it is. In 
these respects gas evolution is completely analogous to precipita- 
tion. Since all gases are less soluble at high than at low tempera- 
tures, gas evolution is always more complete the higher the 



294 Introduction to General Chemistry 

temperature. Gas evolution and precipitation differ in one 
very important additional respect: at a given temperature the 
M.S. of a precipitate has a fixed value, while that of a gas depends 
upon the pressure of the gas above its solution. In most cases 
the total solubility and therefore also M.S. of a gas is directly 
proportional to its (partial) pressure (Henry's Law, 126). If 
during gas evolution the partial pressure of the gas given off 
is kept down by removing the gas (as by blowing it away with a 
stream of air) as fast as it is liberated, its M.S. will be reduced to 
a vanishingly small value. In consequence the dissolved gas 
will be practically or even completely removed from the solution. 
Thus, in the reaction between NaOH and NH 4 C1, if a current of 
air is blown through the solution every trace of NH 3 will finally 
be removed, so that the reaction will be absolutely complete. 
The remaining solution will contain only common salt. The 
same result is attained if the solution is boiled, in which case the 
evolved steam takes the place of the air current. The high 
temperature also hastens the removal of the NH 3 . All reactions 
giving gases which follow Henry's Law may be driven to com- 
pletion by the continuous removal of the gas by means of a 
current of an inert gas or by steam produced when the solution 
is boiled. We have seen (455, 461) that the reason why a little 
soluble salt dissolves is the efficient removal from solution of one 
or both of its ions to form some new substance, which of course 
must be soluble or volatile if the resulting mixture is to be a 
solution. In the reactions studied so far the removal of ions 
has been accomplished by the formation of little ionized or 
little soluble substances. There are other ways of removing 
ions. These we shall take up later. We shall find that the 
dissolving of little soluble substances in question depend upon 
the same fundamental principle, and that these new cases differ 
from those studied in this chapter only in secondary ways, the 
means by which the ions in question are removed from the solu- 
tion (532, 560, 626). 

464. The Value of the Ionic Hypothesis. — In chapters xvii 
and xviii, we have applied the ionic hypothesis to the interpreta- 
tion of reactions between acids, bases, and salts and have seen 



Applications of the Ionic Hypothesis 295 

that this hypothesis leads to fairly simple explanations of a 
great variety of facts. Furthermore we have seen that if we 
know the degrees of ionization and the solubilities of the sub- 
stances concerned in any reaction we are able to predict what 
the result of the reaction will be. Herein lies the enormous 
practical value of the ionic hypothesis. 

In chapter xvii (412) we called attention to some of the 
glaring inconsistencies of the hypothesis; but we have also 
pointed out that the practical value of any hypothesis is not its 
truth but its usefulness. Having now, we hope, shown its use- 
fulness, we shall in later chapters consider whether it is true 
(chaps, xx, xxvii). 



CHAPTER XX 
ELECTROCHEMISTRY 

465. Introduction. — The present chapter will deal first with 
some of the marvelous developments of our knowledge of elec- 
tricity and matter during the last two decades. We now have 
good reason for believing that electricity like matter is of a 
granular or " atomic" nature. The grains or "atoms" of free 
electricity are all exactly alike and of the variety known as 
negative electricity. These grains are called electrons. Posi- 
tive electricity is not known in a free state; that is, it is only 
known as a positive charge on ordinary chemical atoms or 
larger masses of matter. 

466. The Granular Nature of Electricity; Electrons. — In 
chapter xvii (403) it was shown that Faraday's Law of Electro- 
chemical Equivalents leads directly to the conclusion that all 
univalent ions, in solution, carry equal charges of electricity (404) . 
The charge on a single univalent ion may be called a unit charge. 
Each bivalent ion has two unit charges, each trivalent ion three 
unit charges, etc. As early as 1874 Storey pointed out that these 
facts indicate that electricity is granular in nature, that each 
univalent atom is associated with one such granule to form a 
univalent ion, that a bivalent ion is an atom with two granules 
of electricity, etc. Furthermore a little later he proposed to 
call the quantity of electricity of a single granule an electron. 
This quantity is exceedingly minute. The common unit of 
quantity, one coulomb, is equal to more than a billion billion 
electrons. According to present-day usage the term electron 
means a single electronegative granule of electricity. 

467. Proof of the Existence of Electrons. — Although evidence 
was gradually accumulating during the last quarter of the nine- 
teenth century, it was not until more recently that positive 
evidence was obtained that electricity is granular and is made 
up of electrons. The crowning work was that of Professor 

296 



Electrochemistry 



297 



Robert Millikan, an American physicist who showed that when 
a very small sphere is charged with more and more electricity 
the quantity of electricity increases by small, equal additions, and 
not continuously. This is exactly what we should expect if the 
charge is made up of a small number of electrons. 

The spheres used were oil drops of microscopic size, not 
much larger, in fact, than the particles of dust that can be seen 
floating in the air when a beam of sunlight penetrates a dark- 
ened room. A drop was made visible by a beam of bright light 
and was viewed through a short-focus telescope. In still, dust- 
free air the drop fell, under the action of gravity, at a constant 
velocity that could be. accurately measured. It is interesting 





•J 
















3 


M 


+ 








•fir 




•0 








~r 


1 — 1 


c 


()L* 


T 




N 




V. 


5 








1 E 





Fig. 75 



to note that although every body, however small, will fall with 
steadily increasing velocity in a vacuum, a very small body 
falls with constant velocity in air, owing to the viscosity of the 
latter. The drops used fell with a velocity of about one milli- 
meter per second. The principal parts of Millikan's apparatus 
are shown in Fig. 75 : M and N are parallel metal plates insulated 
from one another and connected through a switch to the terminals 
of a high-potential battery, B. The upper plate, M, can be 
charged positively and the lower one, N, negatively. A minute 
oil drop, D, is caused to fall into the space between M and N 
through a pinhole in the center of M , and its rate of fall is 
measured while M and N are uncharged, observations being 
made with a telescope, T. A minute negative charge is then 
given to the drop (in a way that need not be considered at 
present), and the plates M and N are charged. The drop is 



298 Introduction to General Chemistry 

now attracted by M and repelled by N, so that it moves upward. 
When it is close to M the electric field is switched off (S, switch; 
E, earth), and the drop is again allowed to fall, and its speed 
(time of fall) is again measured. With uncharged plates (field 
off) the drop falls at exactly the same rate, whether it is charged 
or uncharged. Next its speed upward is measured with the 
field on. This speed is always the same as long as the charges 
on M and N remain constant (constant field), and the charge 
on the drop is unchanged. But increase of negative charge on 
the drop increases the upward speed, and decrease of negative 
charge decreases the upward speed, the field remaining constant. 
The speed upward is a measure of the force of electrical attrac- 
tion by M and repulsion by N of the charged drop and is there- 
fore a measure of the charge on the drop. A drop could be made 
to make hundreds of trips up and down. The downward 
velocity (field off) was always the same; the upward velocity 
(field on) varied with the charge. The charge for each upward 
speed was found by a simple calculation. 1 It turned out that 
the charge on the drop was in every case a multiple by a whole 
number of the smallest possible charge on the drop. Thousands 
of observations were made in these experiments, but not an 
exception was found to the foregoing statement. This proves 
conclusively that electricity is granular in nature. It has been 
shown in other ways that the granules of electricity composing 
the charge on an oil drop are of the same magnitude as the unit 
charges of ions of electrolytes in solution. We may therefore call 
them electrons and say that all electric charges are made up of 
one or more electrons. In Millikan's experiments the oil drops 
used were observed to carry all possible charges from a single 
electron to over a hundred electrons; in no single case was a 
fraction of an electron found. The electron is therefore the 
smallest indivisible particle of electricity. 

468. The Nature of an Electric Current. — The relation of 
an electric charge to an electric current was first clearly estab- 
lished in 1876 by the American physicist Rowland, who showed 

J A popular account of Professor Millikan's work is given in his book, The 
Electron. Chicago: The University of Chicago Press, 191 7. 



Electrochemistry - 299 

that when an electrically charged gilt disk was very rapidly 
rotated it produced the same sort of deflection on a magnetic 
needle as that due to a current of electricity flowing through a 
wire having the same position with respect to the needle as 
that occupied by the disk, Fig. 76. This experiment proved 
that a current of electricity is nothing but an electric charge 
in motion, just as a current of water is nothing but water in 
motion. 

469. The Electron Theory of Electric Currents. — If we 
accept the view that an electric current is an electric charge in 
motion, and also take into account the fact that an electric 
charge is merely an assemblage of electrons, we are at once led 
to the supposition that a current in a wire 
is only a stream of electrons passing through 
the wire. 

470. The Structure of an Atom. — If 
we think of a metal wire as made up of 
" solid," impenetrable atoms, it is not 
very reasonable to imagine that par- 
ticles of electricity (electrons) could pass 
through it. However, physicists and FlG 6 
chemists have in recent times come to 

the conclusion that an atom is by no means a homogeneous, 
solid lump, but that it is a rather complex structure, con- 
sisting largely or even wholly of negative electrons rotating 
in more or less circular orbits about a positively charged 
nucleus. The sum of the negative charges of the electrons 
is exactly equal to the positive charge of the nucleus, so that, 
as a whole, an atom has no excess of either kind of elec- 
tricity. The structure of an atom may be likened to that of 
the solar system, in which the sun corresponds to the nucleus 
and the planets to the surrounding electrons. The distances 
between the electrons composing an atom are probably large 
compared with the size of an electron, so that a stray electron 
might pass through an atom with the same ease that a comet 
passes through the solar system, or a bullet may pass through 
a flock of birds without striking any one of them. 




300 Introduction to General Chemistry 

471. How a Wire "Carries a Current." — If we think of a 
wire as made up of atoms of the sort here pictured, it is easy to 
see how a stream of electrons might pass through it. In a wire 
(not connected with any electrical source) some electrons are 
continuously becoming detached from their original atoms; 
these probably move through and among the atoms, occasionally 
replacing, for the time being, those that have been lost by other 
atoms. A metal always contains more or less of these free, 
wandering electrons, as well as an equal number of atoms which 
are deficient in electrons. When the terminals of a battery are 
joined by a wire, the positive pole of the battery attracts and 
the negative pole repels the free electrons of the wire. This 
causes a drift of electrons along the wire, and this drift of electrons 
constitutes the current in the wire. The progress of electrons in 
the direction of the drift is slow, a matter of a few centimeters 
per minute. The well-known fact that the effect of closing an 
electric circuit is felt almost instantaneously at a great distance 
(as illustrated, for example, by our everyday telephone experi- 
ence) is explained by the assumption that all the mobile electrons 
in the wires of the circuit move forward at the instant the circuit 
is closed. The case is just like that of drawing water from a 
supply pipe; water flows out the instant the faucet is opened, 
being pushed forward along the whole pipe by the water forced 
into the pipe by the pump. 

472. The Direction of an Electric Current. — Before the 
nature of an electric current had been discovered it was cus- 
tomary to consider that the current in the wire flowed from the 
positive to the negative pole. Since the drift or flow of electrons 
is in the opposite direction, there is danger of misunderstanding 
in speaking of the direction of the current. It is perhaps best 
to speak of the direction of the negative current, which is then 
the direction of drift of the electrons. 

473. Nonconductors of Electricity. — All metals are good con- 
ductors, but the non-metals are practically nonconductors or in- 
sulators. To .account for this difference we have only to suppose 
that a non-metal, like sulfur, contains but very few free or mobile 
electrons and therefore has very little ability to carry a current. 



Electrochemistry 301 

474. Production of Electric Charges by Friction. — If a glass 
rod is rubbed with a piece of silk, the former takes on a positive, 
the latter a negative, charge. This is explained by assuming 
that a few stray electrons of the glass have been " wiped off" by 
the silk. The rubbing of the glass by the silk is of importance 
only in insuring intimate contact between the two. Another 
example of similar nature is found in the familiar electrification 
of the hair when combed with a hard-rubber comb in dry weather. 
Here the comb acquires a negative charge and the hair a posi- 
tive one. In general, when any two different substances are brought 
together they become electrified with opposite charges. This may be 
taken to mean that electrons accumulate in excess more easily 
on some kinds of matter than on others. 

475. Cathode Rays.— \ 

When a high- voltage electric r^± ^ : ±L- :"■ ■'/- I'j- 1 'irlM ~ " '•'"'""" ' ' rill \ 

current is passed through a ^~^*~ - -" : ^-= -^ =3t r.'?:; ;:;; jjif / 

Crookes tube, Fig. 77, which " ""A* * A-L-____^/ 

is an evacuated glass bulb \J 

< » 

having a metallic cathode, FlG 

C, and an anode, A, rays, 

known as cathode rays, are given off by the cathode and cause 
a yellowish-green fluorescence of the opposite end of the tube. 
These rays are readily stopped by a sheet of metal, as shown by 
the fact that a screen (in the form of a Maltese cross) casts its 
shadow on the glass. Even transparent substances like glass 
do not transmit the cathode rays any better than do metal sheets 
of comparable thickness. Extremely thin sheets of material like 
aluminum or gold leaf permit partial transmission of the cathode 
rays. 

476. X-Rays.^-Cathode rays produce X-rays, also known as 
Roentgen rays, which radiate from any object struck by the 
former. A modern X-ray tube is shown in Fig. 78. This is a 
modified Crookes tube, intended for the use of large currents 
and the production of powerful X-rays. The cathode rays come 
from the specially constructed cathode, C, and strike a target, T, 
made of metallic tungsten, which metal is chosen because of its 
very high melting-point (3000 ). When the cathode rays are 



302 Introduction to General Chemistry 

stopped by the target, part of their energy is transformed into 
X-rays, and the balance appears as heat, so that the target 
becomes red, or even white, hot. Recent work has proved that 
the X-rays, which are very different from the cathode rays, are, 
like visible light, vibrations of the so-called luminous ether and 
differ from visible light in having wave-lengths about one- 
thousandth as great as the latter. 

477. The Nature of Cathode Rays. — The extensive investiga- 
tions of Sir William Crookes on cathode rays, during the seventies 
of the last century, led this famous English physicist -to con- 
clude that these rays were matter in a highly rarefied or 




Fig. 78 

ultra-gaseous state, which he called a fourth state of matter 
(the other three states being solid, liquid, and gaseous). And 
in the light of our present knowledge of the real nature of these 
remarkable rays we must admit that Crookes 's conclusion was 
substantially correct, although it was by no means the last word 
on the subject. 

It has long been known that cathode rays travel in a straight 
line in a vacuum, but that they may be deflected in an arc of a 
circle by a transverse magnetic field. The apparatus shown in 
Fig. 79 serves for lecture demonstration of this interesting 
phenomenon. A narrow beam of rays coming from the cathode 
and passing through a slit in a mica plate strikes along a screen 
covered with a specially prepared form of zinc sulfide, which 
becomes luminous in the line where it is struck by the rays. If 
now a horseshoe magnet is presented so that the N pole is above 



Electrochemistry 



303 



the plane of the paper and the S pole below it, the beam is 
deflected to the position of the curved line. 

It is a well-known fact that a wire, free to move, is deflected 
by a magnetic field when a current is passed through it. The 
direction of deflection of the wire is determined by the direction 
of the current in the wire. The deflection of the cathode rays 
by a magnetic field indicates that the rays are electricity in 
motion, the direction of deflection corresponding to that of a 
stream of negative electricity coming from the cathode, which is, 
of course, the negative electrode. If we grant that the current 
in the wire leading to the cathode is, in reality, only a stream of 
negative electrons in the wire, we have only to suppose that these 




electrons do not stop on reaching the cathode, but shoot out from the 
surface of the latter and thus constitute the cathode rays. 

478. Proof that the Cathode Rays Are a Stream of Electrons. 
— The conclusive proof that the cathode rays are a stream of 
negative electricity (presumably electrons, since all negative 
charges consist of electrons) was furnished by the work of 
Perrin, a French physicist. Perrin's apparatus is shown in 
Fig. 80. It was a special form of Crookes tube having the 
cathode at C, the anode at A, and at B an insulated metal 
receiver, into which the cathode rays could be deflected by means 
of a magnet. This receiver was connected by a wire to an elec- 
troscope, capable of detecting any electric charge given to the 
box and determining its sign, whether positive or negative. 
When the cathode rays were started no charge passed into the 
receiver until the rays were magnetically deflected so as to fall 
into the receiver; then the latter acquired a large negative 



304 Introduction to General Chemistry 

charge. To guard against stray electric charges the receiver 
was surrounded by a metal shield connected to the earth, E. 
The experiments above described, together with many other 
facts, have led to the conclusion that the cathode rays are com- 
posed of negative electrons shot out from the cathode with high 
velocity. 

479. The Mass of an Electron. — An electron behaves as 
thought it had mass. In the first place we know that moving 
electrons have energy, since the cathode rays can produce light, 
heat, and X-rays, all of which are forms of energy. Since the 
kinetic energy of a moving body is proportional to the product 



Fig. 80 

of its mass and the square of its velocity, we can account for 
the energy of the cathode rays by assuming the electrons to 
have mass. Furthermore the fact that it requires an appreciable 
magnetic force to deflect the cathode rays and that the extent 
of the deflection (for rays of a given velocity) is proportional 
to the strength of the magnetic force is also evidence that elec- 
trons have mass. One of Newton's laws is to the effect that a 
moving mass continues in a straight line unless acted upon by a 
transverse force. Conversely, if a force is required to deflect 
a moving electron, we are warranted in assuming that the latter 
has mass. By methods that we cannot explain here it has been 
shown that the mass of an electron is about one eighteen-hundredth 
that of an atom of hydrogen. 

480. The Beta Rays of Radium. — The spectacular properties 
of radium have been brought to the attention of nearly everyone, 



tdectro chemistry 305 

whether he is a student of chemistry or not. Radium gives out 
three kinds of rays, the alpha, beta, and gamma rays. Of these 
the beta rays very closely resemble the cathode rays. Like 
cathode rays they are deflected by a magnetic field in a direc- 
tion which indicates that they too are a stream of electrons 
shot out with high velocity from the radium. Radium is, by 
all ordinary tests, an element. It resembles barium as closely 
as potassium resembles sodium. Here then is an element that 
spontaneously gives off negative electricity in the form of elec- 
trons shot out with great velocity. 

The alpha rays have been proved to be atoms of the element 
helium, He (atomic weight = 4), each of which carries a double 
positive charge. These rays are also shot out with high velocity. 
The gamma rays are identical with X-rays. 

481. The Disintegration Hypothesis. — The extraordinary 
behavior of radium has been satisfactorily explained by the 
disintegration hypothesis of Rutherford and Soddy. These 
scientists assumed that a radium atom is not a homogeneous 
solid particle but a very complex structure made up of electrons 
revolving rapidly in more or less circular orbits about a nucleus 
of positive electricity in the manner already described. It is 
further assumed that an atom of radium may become unstable 
and throw of a single electron {beta ray) or a larger mass {an atom 
of helium, which is an alpha ray), leaving behind an atomic residue 
of smaller mass and therefore smaller atomic weight. This hy- 
pothesis is in complete accord with all known facts concerning 
radium and radioactive phenomena. 

482. The Electrical Nature of Matter. — The study of radio- 
active substances, of which, in addition to radium, about thirty 
are known, has led to the conclusion that the atoms of all ele- 
ments, whether radioactive or not, are constructed on the same 
same plan as that of radium*. According to this hypothesis the 
atom of one element differs from that of another only in the number 
and arrangement of the electrons composing it. The mass of an 
atom is, at least in part, accounted for by the mass of the elec- 
trons composing it. All matter is considered to be of electrical 
origin. 



306 Introduction to General Chemistry 

483. The Nature of an Ion. — A single sodium ion is an atom 
of sodium having a single positive charge of electricity equal in 
quantity but opposite in sign to that of an electron. The 
simplest explanation of the difference between an ion and an 
atom of sodium is found in the assumption that the ion is an 
atom which has lost one electron. The atom was originally elec- 
trically neutral, because the positive charge of its nucleus was 
just equal to the sum of the negative charges of its surrounding 
electrons. If one electron is lost, the atom will have an excess 
positive charge just equal in magnitude to that of one electron. 
Since metallic atoms all form positive ions we conclude that all 
such atoms are able to lose electrons. Moreover, an atom of 
a univalent metal can lose but one electron and its ion will have 
a single unit charge, thus, 

Na(atom)->Na + +one electron. 

A bivalent atom can lose two electrons, 

Ca(atom)->Ca ++ +two electrons. 

A trivalent atom, such as that of aluminum, can lose three elec- 
trons, etc. 

Later work has shown that the ions are undoubtedly hydrated 
to some extent. The actual formula of sodium ion might be 

represented thus: 

Na(H 2 0),+ . 

The subscript x represents a small integer, probably 2 or 3. In 
practice we do not include the water in the formula, since in 
the first place the exact data necessary are wanting, and in the 
second place the relationships in our reactions seem to be satis- 
factorily represented without it. 

484. Valence. — The idea just presented leads to a simple 
explanation of valence (147, 183). The metals which form 
only positive ions do so by the loss of one or more electrons from 
each atom. The valence of an atom of a metallic element is 
determined by the number of electrons it has lost. 



Electrochemistry 



307 



A negative ion, such as CI", is an atom which has taken up 
an extra electron. Atoms of metals do not take up extra elec- 
trons. Only the atoms of non-metallic elements behave in this 
way. The valence of a negative ion consisting of one atom corre- 
sponds to the number of electrons the atom has acquired. 

485. Theory of the Union of Sodium and Chlorine. — It is 
well known that sodium and chlorine unite very energetically 
to form NaCl. The simplest explanation of the cause of union 
is found in the assumption that an atom of sodium has a great 
tendency to lose an electron, and that an atom of chlorine has a 
great tendency to take up an extra electron. The violent reac- 
tion that we observe when we bring these two elements together 
is only the outward manifestation of the passage of electrons 
from the atoms of sodium into the atoms of chlorine. The 
residue of the sodium atom now has an excess of positive elec- 
tricity, while the chlorine atom with its extra electron is charged 
negatively. Since unlike electric charges attract each other, 
we may well assume that the two parts of the NaCl molecule 
are held together by electrical attraction. 

486. The Cause of Ionization. — If two insulated bodies are 
oppositely charged, Fig. 81, the force with which they attract 




Fig. 81 



each other is proportional to the product of their charges and 
inversely proportional to the square of the distance between 
them. There is, however, one additional factor that determines 
the strength of the attraction, and that is the nature of the 
surrounding medium. Usually this is air. If the medium were 



308 Introduction to General Chemistry 

glass instead of air the attraction would be only about one-third 
as great, other things remaining the same; but if the medium 
were water the attraction would be only one-eightieth as great 
as for air. If then we dissolve NaCl in water, the molecules are 
surrounded by a medium which lessens enormously the attract- 
ive force which holds their parts together; as the result, mole- 
cules will tend to fall apart, thus, 

NaCl->Na++Cl-. 

The positive part is the sodium ion, the negative part the chlorine 
ion. According to this explanation the molecule of salt before 
it ionizes is made up of two oppositely charged parts. These 
are not ordinary atoms, since the one has lost an electron which 
the other has gained. We ought to say that the NaCl mole- 
cule is made up of a sodium ion (Na + ) electrically bound to an 
ion of chlorine (Cl~). The act of ionization, which takes place 
when the salt is dissolved, is only the falling apart of the ions already 
present, on account of the great decrease in attractive force caused 
by the surrounding water. In other words, molecules are com- 
posed of bound ions, while in solution part of the ions are free. 
The ionization of all acids, bases, and salts is explained in pre- 
cisely analogous fashion to that in the case of NaCl. 

487. The Electronic Description of Electrolysis. — According 
to the electronic description of electrolysis, when an ion reaches 
an electrode it either gains electrons or loses them. Thus the 
positive ions Cu ++ , Pb ++ , H + , etc., each gain enough electrons 
to make them electrically neutral ; while Cl~, I~, S~~, and 
0~~ each lose electrons and become free elem'ents. 

488. The Displacement of Non-metals by One Another. — It 
will be recalled that chlorine acts on a solution of hydrobrorhic 
acid or any bromide, setting free bromine, thus: 

Cl a +2HBr->Br a +2HCl. (259) 

Similarly bromine acts on iodides, as, for example : 
Br 2 +2KI->I 2 +2KBr. 



Electrochemistry 309 

Iodine acts on H 2 S, in solution, setting free sulfur: 

I 2 +H 2 S->S+2HI. (339) 

The order in which the four above-mentioned elements displace 
one another is therefore as follows : 

CI, Br, I, S. 

Each will set free from its compounds any one following it. We 
may also include fluorine and oxygen in the series, and, since 
fluorine will displace any of the other elements mentioned, it will 
head the series. The position of oxygen is determined by the 
fact that a H 2 S solution reacts with atmospheric oxygen to form 
free sulfur and water, 

2 + 2 H 2 S-> 2 S+2H 2 0, (339) 

and that a solution of HI also reacts with oxygen of the air 
(slowly) to form water and free iodine, 

2 + 4 HI^2l 2 + 2 H 2 0. (265) 

On the other hand, HBr solution is scarcely affected by oxygen 
gas, and HC1 solution not at all. Oxygen will therefore precede 
iodine and sulfur and follow bromine in the list. The whole 
displacement series is then as follows: 

F, CI, Br, O, I, S. 

489. Electronic Interpretation of Displacement. — If the re- 
action 

Cl 2 +2HBr^Br 2 + 2 HCl • 

takes place in very dilute solution, the two acids are nearly com- 
pletely ionized, and we may leave the H + ion out of considera- 
tion. The reaction in its simplest aspect is as follows: 

Cl 2 + 2 Br-^Br 2 +2Cl-. 

This means that each Br~ ion loses an electron, which, passing 
into a chlorine atom, changes the latter into a Cl~ ion. We 
conclude that chlorine atoms take up electrons more readily than 



310 Introduction to General Chemistry 

do atoms of bromine. Considering next the six elements of the 
displacement series, we may say that fluorine has the greatest 
tendency to take up electrons, and sulfur the least, and that the 
tendencies of the other elements come in the order indicated in 
the list as given. Summarizing, we may say that of two elements 
of the above-mentioned series the one whose atoms have the greater 
tendency to take up electrons will set the other free from its compounds 
with positive ions. 

490. The Displacement of Metals by One Another. — Strips 
of metallic zinc placed in solutions of salts of lead, copper, 
mercury, and silver will react as indicated by the following 

equations : 

Zn+Pb(N0 3 ) 2 ->Pb+Zn(N0 3 ) 2 , 
Zn+CuS0 4 ->Cu+ZnS0 4 , 
Zn+HgCl 2 -»Hg+ZnCl 2 , 
Zn+2AgN0 3 ->2Ag+Zn(N0 3 ) 2 . 

In other words, zinc displaces each of the above-mentioned metals 
from its salts. 

If strips of metallic lead are placed in solutions of salts of 
zinc, copper, mercury, and silver, no reaction takes place with 
the zinc salt; but the other three metals are set free, while the 
lead atoms pass into solution as positive ions. In similar 
fashion metallic copper sets free mercury and silver from their 
salt solutions, but it does not affect solutions of zinc or lead salts. 
Mercury displaces silver from its salts but has no action on salts 
of zinc, lead, or copper. Metallic silver will not displace from 
their salts any of the other metals just considered. The order 
of displacement of the five metals is therefore as follows : 

Zn, Pb, Cu, Hg, Ag. 

491. Electronic Interpretation of Metallic Displacement. — 

The action of zinc on solutions of copper salts may be represented 
in simplified form thus: 

Zn+Cu++->Cu+Zn++. 

This means that an atom of zinc gives up two electrons to an 
atom of copper. Since zinc displaces copper equally well from 



Electrochemistry 311 

solutions of all its simple soluble salts, we conclude that an 
atom of zinc has a greater tendency to lose electrons than has 
an atom of copper, but, since metallic copper displaces silver 
from any of its salts, thus, 

Cu+2Ag+-> 2 Ag+Cu++, 

we conclude that an atom of copper has a greater tendency to 
lose electrons than has an atom of silver. 

The order of the metals in the displacement series 

Zn, Pb, Cir, Hg, Ag 

is therefore the order in which they fall according to the decreasing 
ease with which their atoms tend to lose electrons. In the case of 
any two metals of the preceding list the one whose atoms have 
the greater tendency to lose electrons will set the other free from 
its compounds with negative ions. 

492. A More Complete Displacement Series of Metals. — 
Most of the familiar metals may be included in a single displace- 
ment series, which shows at the same time the tendencies 
of the atoms of the metals to lose electrons and so change 
into positive ions. The list is given in Table XIX. In this 





TABLE XIX 


Displacement Series of the 


Potassium 


Nickel 


Sodium 


Tin 


Barium 


Lead 


Calcium 


Hydrogen 


Magnesium 


Copper 


Aluminum 


Mercury 


Zinc 


Silver 


Iron 


Platinum 


Cobalt 


Gold 



list (in which the second column follows the first) each metal 
tends to displace, or set free from its combination with negative ions, 
any element which follows it. 

Hydrogen has been placed in the list between lead and copper. 
Any metal above hydrogen in the series will react with a normal 



312 Introduction to General Chemistry 

solution of hydrochloric acid to set free hydrogen (at atmospheric 
pressure) and pass into solution as chloride. The metals follow- 
ing hydrogen in the list do not react readily, if at all, with hydro- 
chloric acid. The first four elements of the series react with 
water to set free hydrogen. Therefore metallic potassium 
placed in a solution of NaCl does not set free metallic sodium 
but causes the evolution of hydrogen. The order shown for 
the first four elements of the table is in fact that of their tend- 
encies to lose electrons as determined by means other than 
direct displacement. 

493. The Production of an Electric Current. — In the reac- 
tion between zinc and copper sulfate the essential change, as we 
have seen, is that represented by the equation 

Zn+Cu++->Zn++ + Cu. 

We have said that this change is the result of the passage of 
two electrons from each atom of zinc into an atom of copper. 
Now if this is true we ought to be able to get an available elec- 
tric current from the reaction; but if a piece of 
Zn zinc is dipped into a solution of a copper salt, 

Fig. 82, no evidence of the production of such 
a current is to be observed. How indeed could 
we expect to detect the production of an electric 
current under the conditions pictured in Fig. 82? 
CuSOe If a passage of electrons occurred, it would be 
between the zinc rod and the layer of copper 
FlG g2 ions in the solution surrounding the rod, and we 

could not readily detect such a current, much 
less make any use of it. If we wish to make this supposed 
current available for detection and use, we must so arrange the 
reacting substances that the Cu + . + ions are not directly in con- 
tact with the zinc rod, and then provide a wire for the transfer 
of electrons from the zinc rod to the copper ions. This can be 
done by arranging the four substances of the reaction 

Zn+ CuS0 4 ->ZnS0 4 + Cu 



^mFk± 



Electrochemistry 



3*3 



in the manner shown in Fig. 83. Here we have a zinc rod dipping 
into a solution of ZnS0 4 in one beaker, and a copper rod dipping 
into a solution of CuS0 4 in the other beaker. A glass tube 
filled with ZnS0 4 solution and loosely stoppered with cotton 
plugs forms a so-called salt bridge between the two beakers. If 





<^^^ 


3 












Zn 
ZnSO, 




+ 

|Cu 




'"". 




m 


! ^ 


CuSO, 


r— 



Fig. 83 

now the two rods are connected by wires to a galvanometer, a 
current is found to flow in a direction indicating the passage of 
electrons from the zinc rod through the wire (and galvanometer) 
to the copper rod. At the same time metallic copper begins to 
deposit on the copper plate, and metallic zinc begins to pass into 
solution. In fact, the reaction 

Zn+ CuS0 4 ->ZnS0 4 + Cu 

begins to take place just as soon as the metallic circuit is closed 
between the upper ends of the zinc and copper rods. No action 
occurs before the circuit is closed, and all action stops when the 
circuit is broken. 

494. The Mechanism of Current Production. — In detail the 
actions that occur with closed circuit are as follows: zinc atoms 
pass from the rod into the solution, each atom of zinc leaving 
behind two electrons and changing thereby into a Zn ++ ion. 
.The electrons thus liberated flow through the wire to the copper 
rod in the CuS0 4 solution. Copper ions in contact with the 
copper rod take up two electrons each, being thereby changed 
into ordinary copper atoms. These latter adhere to the copper 
rod as a metallic coating. Fresh Cu ++ ions move up to the 
copper rod by diffusion, so that, as the ions in contact with the 
rod take on electrons and change into copper atoms, others move 
up by reason of their kinetic motion to take their places. On 
the other hand Zn ++ ions, newly formed at the zinc rod, diffuse 



314 Introduction to General Chemistry 

out into the solution. These changes tend to cause a deficiency 
of S0 4 ~~ ions about the zinc rod, and an excess of the same ions 
about the copper rod. The attraction between the excess of 
S0 4 ~~ ions, on the one hand, and the excess of Zn ++ ions, on 
the other, causes a migration of these ions in opposite directions 
through the solution and the salt bridge, and thus serves to 
maintain in each cubic centimeter of the whole solution as many 
negative as positive ions and thereby to keep the solution, as a 
whole, electrically neutral. 

495. The Function of the Salt Bridge. — The necessity of 
some sort of connection between the solutions of ZnS0 4 and 
CuS0 4 in the two beakers, Fig. 83, is obvious. If we remove 
the salt bridge, which in this case is a ZnS0 4 solution, the circuit 
is broken, and all action comes to a stop. By the use of the 
bridge we are able, by placing the CuS0 4 in a separate beaker, 
to keep it from coming in contact with the Zn rod. The use of 
a metal wire in place of the salt bridge would apparently be a 
simpler plan, but it would not serve, because new products 
would be set free by electrolysis at each end of the wire. We 
could, however, use in the bridge, instead of the ZnS0 4 , a solu- 
tion of CuS0 4 or, in fact, of NaCl or almost any other salt. In 
case the bridge contains NaCl, the Na + ions serve in place of 
Zn ++ to carry the positive charge from the ZnS0 4 solution 
to the CuS0 4 solution, and the Cl~ ions to carry the negative 
charge in the opposite direction. 

496. Galvanic Cells. Electric Batteries. — A galvanic cell, 
or, as it is more popularly known, an electric battery, is any 
kind of apparatus by means of which an electric current is 
produced by chemical reactions. Dry batteries and storage bat- 
teries are, at present, the most familiar types. The first prac- 
tical form of the zinc-copper cell just studied was known as 
the Daniell cell; a later modification is known as the gravity 
battery. Properly speaking, the term battery means a group 
of cells, but this term is frequently used at present to mean a 
single cell. 

497. The Gravity Battery. — A gravity cell is shown in 
Fig. 84. It consists of thin sheets of copper surrounded by a 




Electrochemistry 315 

solution of copper sulfate in the lower half of the glass jar and a 

heavy zinc "crowfoot" surrounded by a zinc sulfate solution in 

the upper half. Attached to the copper sheets is an insulated 

copper wire. A new cell is set up by filling the jar with water, 

placing the copper and zinc in position, and adding more than 

enough solid CuS0 4 (blue vitriol or bluestone) to saturate the 

lower layer. No ZnS0 4 need be added; instead, 

twenty or thirty g. of NaCl are sprinkled into 

the water. The solution is not stirred. The 

CuS0 4 gradually dissolves, giving a saturated 

solution which soon fills the lower part of the 

cell. If now the insulated wire leading from 

the copper is connected to the zinc, a current 

flows through the wire, and the changes already 

described take place. The NaCl is used to ^ 

r m Fig. 84 

make the water conduct the current prior 

to the formation of sufficient ZnS0 4 for this purpose. Until 

recently the gravity battery was used to operate all telegraph 

lines. 

498. Other Kinds of Galvanic Cells. — It is possible to make 
a cell that will give a current by the use of any pair of metals 
(not acted upon by water) , each surrounded by a solution of one 
of its salts. In each cell the experimental arrangement may be 
that shown in Fig. 83. 

499. Electromotive Force and Voltage. — A body at rest can 
be set in motion only by the action of a force (Newton's law). 
In a similar manner we assume that the current (stream of 
electrons) produced by a battery is the result of an electrical 
force called the electromotive force, E.M.F. The unit of 
E.M.F. is the volt (named after the pioneer electrical experi- 
menter Volta). The gravity cell has an E.M.F. of 1. 1 volts. 

The farther the two metals forming the electrodes of a cell 
of any kind are removed from each other in the displacement 
series (492) the greater the E.M.F. of the cell. The reason for 
this is found in the fact that the metals heading the list give 
off electrons most readily (with greatest force). The order in 
the list represents, in fact, the relative force with which the 



316 Introduction to General Chemistry 

element loses electrons. The difference of such forces for the 
two metals of a cell is, for practical purposes, the chief deter- 
mining factor of the E.M.F. (voltage) of the cell. This difference 
of forces between the electrodes is also often called the potential 
difference of the electrodes. 

There is another important factor to be considered besides 
the nature of the reactions at the electrodes, and that is the con- 
centration of the ions in solution. For example, the more con- 
centrated the copper ions at the copper electrode the faster is the 
reaction carried on by these ions at a given temperature. Now 
the difference of potential at the terminals of a cell is a measure 
of the rate of the reaction in progress ; hence it will be increased 
or decreased by concentration changes in the solutions. To 
make careful comparison of the electromotive forces of cells the 
concentrations of the ions must therefore be taken strictly into 
account. However, in the series we are considering, no moderate 
variation of the concentrations of the ions from those found in 
the ordinary laboratory reagents (o.oi to 6N approximately) 
will produce results different from those described here, in the 
•cases under consideration. The effect of the concentration of 
ions on cell potentials should be considered in an exact study of 
the latter subject. 

500. Electrical Energy. — Electrical energy always depends 
on two factors, voltage and quantity of electricity. The unit 
of electrical energy is the joule, named after J. P. Joule, the 
celebrated English scientist, whose work on the mechanical 
equivalent of heat was discussed earlier (370). One joule is the 
amount of energy produced when a quantity of one coulomb of elec- 
tricity flows through a conductor, the ends of which have a potential 
difference of one volt. In general, joules = voltsX coulombs. 
For example, if a gravity cell of 1 . 1 volts E.M.F. delivers 10 
coulombs, the electrical energy produced is 1.1X10 = 11 joules. 
Since the joule is an energy unit, its value is expressible in other 
energy units. Careful experiment has shown that 

1 joule = 10,200 g.cm., 
1 joule = 0.24 calorie, 
1 calorie = 4. 1 8 joules. 



Electrochemistry 



3i7 



It is electrical energy which a consumer pays for and uses. 
The same number of electrons go back to the positive pole of a 
battery as leave the negative pole, but they lose energy in so 
going. The energy which the electrons give up may be liberated 
as heat or may be converted into work by means of devices like 
the motor. 

501. Electronic Explanation of Oxidation-Reduction Reac- 
tions.— The action of chlorine on ferrous chloride in solution 
(i73> 33 2 ) is a simple, typical example of an oxidation-reduction 
reaction, 

2FeCl 2 +Cl 2 ->2FeCl 3 . 

This reaction in dilute solution may be represented by the simpli- 
fied equation 

2Fe++ + Cl 2 -»2Fe+++ + 2Cl-. 

The ferrous ion, Fe ++ , which is the reducing agent, is oxidized 
to Fe +++ by the chlorine atom, which is the oxidizing agent. 
This is explained by assuming that the Fe ++ ion (which is an 
iron atom that has already lost two electrons) gives up a third 
electron, which, passing into the CI atom, changes the latter 
into a Cl~ ion. Thus we see that the oxidation of the Fe ++ ion 




Fig. 85 

consists in its loss of an electron; and the reduction of the CI atom 
consists in its gain of an electron. 

502. Oxidation-Reduction Cells. — The transfer of electrons 
which occurs in the reaction just studied can be made to yield 
ah available electric current quite as readily as that which 
takes place in the reaction between metallic zinc and copper 
sulfate (493). We may demonstrate this fact by means of the 
arrangement shown in Fig. 85. Platinum electrodes are placed 



318 Introduction to General Chemistry 

in each of two beakers, one of which contains the FeCl 2 solution, 
the other the Cl 2 solution (together with some FeCl 3 or NaCl to 
make the solution conduct). A salt bridge joins the two solu- 
tions. Wires from the electrodes are connected with a galvano- 
nometer, which shows the passage of a current in a direction 
indicating a flow of electrons in the wire from the electrode in the 
FeCl 2 solution to that in the Cl 2 solution. The platinum elec- 
trodes serve as carriers of electrons into and out of the solutions. 
Platinum is superior to any other metal except gold for this 
purpose, because of its very slight tendency to pass into solution 
as ions. 

503. Further Examples of Oxidation-Reduction Reactions. — 
Oxidation-reduction reactions are very common. They may 
all be interpreted in terms of electron transfers, as the following 
additional examples will illustrate. Ferric sulfate is reduced by 
zinc according to the equation 

Fe 2 (S0 4 ) 3 -r-Zn->2FeS0 4 +ZnS0 4 . (335) 

The simplified ionic equation is 

2Fe++++Zn->2Fe+++Zn++. 

Each atom of zinc loses two electrons and changes into a Zn ++ 
ion; these two electrons are taken up, one by each Fe +++ ion, 
which is thereby changed to a Fe ++ ion. The zinc, which loses 
electrons, is the reducing agent and is oxidized by ferric ions, 
which gain electrons and are thereby reduced to ferrous ions. 

The action of ferric salts and soluble iodides is illustrated by 
the following reaction: 

2FeCl 3 +2KI->2FeCl 2 +2KCl+I 2 , 
or in simplified form by 



Fe+++ + 2l-->2Fe++-r-I 2 . 



A closely analogous reaction takes place in the reduction of 
ferric salts by hydrogen sulfide: 

2FeCl 3 +H 2 S^2FeCl 2 +2HCl+S. 



Electrochemistry 319 

The simplified equation is 

2Fe+++ + S~ ->2Fe++ + S. 

The electronic explanation of each of the foregoing reactions can 
easily be made by the student. 

Other more intricate oxidation-reduction reactions, which 
will require a somewhat more extended discussion, will be taken 
up in subsequent chapters. 

In all oxidation-reduction reactions transfers of electrons 
occur; and in all cases the atom or ion which is oxidized loses 
one or more electrons, and the atom or ion which is reduced gains 
one or more electrons. If an ion does not change its charge or 
its composition in the course of a reaction it is neither oxidized nor 
reduced. 

504. The Oxidation and Reduction of Metals. — When a 
metal passes into solution its atoms take on positive charges. 
This means that each atom of a metal loses one or more elec- 
trons when it changes into an ion. Since we have defined 
oxidation as the loss of electrons (501) we can therefore say 
that when a metal changes into its ions it is oxidized. For 
example, in the reaction 

Fe+ CuS0 4 ->FeS0 4 + Cu, 
which we may write in simplified form thus, 

Fe+Cu++->Fe++ + Cu, 

we say that the metallic iron is oxidized to ferrous ions, and the 
copper ions are reduced to metallic copper. We have already 
seen that the further oxidation of Fe ++ to Fe +++ involves a 
loss of one additional electron. 

505. The Oxidation and Reduction of Non-metals. — When 
a non-metal (chlorine for example) passes into solution, its atoms 
take on electrons. We say, therefore, that in such a case the 
element is reduced. Conversely we say that its ions are oxidized 
when by loss of electrons they are changed to atoms of the 
element. 



320 



Introduction to General Chemistry 



506. Oxidation-Reduction Potentials. — Every oxidation- 
reduction reaction can by suitable arrangement be made to 
yield an electric current. The E.M.F. {voltage) of an oxidation- 
reduction cell is the measure of the force with which the reaction tends 
to take place. The stronger the oxidizing tendency of the oxidiz- 
ing agent and the stronger the reducing tendency of the reducing 
reagent the greater the E.M.F. of the cell. A systematic study 
of such cells has shown that all oxidizing and reducing agents may 
be arranged in a series in the order of their decreasing oxidizing 
tendencies and increasing reducing tendencies. 

507. Oxidation and Reduction by Means of the Electric 
Current. — We have shown that oxidation and reduction are 
capable of producing electric currents. There now remains 
to show that an electric current can accomplish oxidation and 
reduction. Two beakers, Fig. 86, are fitted with platinum 



FeCU 




fit 



4 



Fig. 86 



electrodes and joined with a salt bridge, and in one is placed a 
solution of FeCl 3 , and in the other HC1. Upon passing a current 
from a battery of two dry cells so connected that the electrode 
in the FeCl 3 will be the cathode, it will be found that the FeCl 3 
is reduced to FeCl 2 , while at the same time HO is oxidized to 
free chlorine at the anode. The explanation is as follows: The 
battery sends a steady stream of electrons through the wire to 
the cathode; one electron passes from the latter into each 
Fe +++ ion coming in contact with it, changing the Fe +++ 
into Fe ++ . At the anode Cl~ ions coming in contact with 
this electrode give up to it their electrons and change thereby 
into ordinary CI atoms. The latter then unite in pairs to form 
molecules, aggregates of which soon form bubbles that escape 
into the air. In the solution Fe +++ ions are attracted by and 
migrate toward the cathode, while Cl~ ions are attracted by and 



Electrochemistry 321 

migrate toward the anode. Thus the transfer of electricity from 
one electrode to the other in the solution is accomplished by means 
of the moving ions, while in the wire we have a stream of electrons 
set in motion by the battery. A great variety of other oxidations 
and reductions in solution can be accomplished by means of the 
electric current. In fact, since we may consider the change of 
a metal into its ions as an oxidation of the former, and the reverse 
change a reduction of the ions, we may go farther and say that 
all processes of electrolysis result in oxidation and reduction. The 
anode is the seat of oxidation, since it takes up electrons; the 
cathode is the seat of reduction, because it furnishes electrons. 
These statements apply to all electrolyses irrespective of whether 
the substances formed or liberated at the electrodes are elements 
or compounds. 

508. The Conversion of Chemical Energy into Electrical 
Energy. — The production of heat by a chemical reaction has 
been explained (373) as being due to the conversion of chemical 
energy of the reacting substances into heat energy. If metallic 
zinc is placed in a solution of copper sulfate so that the reaction 

Zn+ CuS0 4 ->ZnS0 4 + Cu 

takes place without the production of an available electric 
current, the quantity of heat liberated is 50,100 calories for one 
symbol weight of zinc. If the same amount of zinc reacts with 
copper sulfate in a gravity cell, 2X96,500 coulombs of elec- 
tricity are delivered into the circuit at an E.M.F. of 1 . 09 volts. 
The electrical energy produced is 2X96,500X1.09 = 210,400 
joules. Since 1 calorie = 4. 18 joules, 210,400 joules = 210,400 
-7-4.18 = 50,300 calories. Thus we see that electrical energy 
equivalent to 50,300 calories is produced in a cell, instead of 
50,100 calories of heat produced when the same amounts of the 
substances react directly, without the production of a current. 
The small excess of energy produced in a cell is accounted for by 
the fact, established by experiment, that this amount of energy 
is absorbed as heat from the surroundings as the cell operates. 
Similar results are observed in the energy production of all other 
galvanic cells. The electrical energy produced by any cell is 



322 Introduction to General Chemistry 

equal to the chemical energy liberated or lost, plus or minus an 
additional amount of energy — plus if heat is taken up from the 
surroundings and minus if it is given out to the surroundings. We 
may consider a galvanic cell or battery merely as a device for 
converting chemical energy into electrical energy. 

509. Conversion of Electrical Energy into Chemical Energy. — 
In all processes of electrolysis electrical energy is used up in the 
production of new chemical substances, and we may conclude 
at once that the electrical energy used is changed to chemical 
energy. 



LBJi/19 



